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WORKS TRANSLATED BY
WILLIAM T. HALL

PUBLISHED BY
JOHN WILEY & SONS, Inc.

F. P. TREADWELL’S ANALYTICAL CHEMISTRY
In Two VOLUMES
Vol. I. Qualitative Analysis.
8vo, xiii +538 pages. Cloth, $3.00

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A. Classen’s Quantitative Analysis by Electrolysis.

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' 410

ANALYTICAL CHEMISTRY

Based on the German text
OF

FB. P. ee Vat, Pu. D.,

Professor of Analytica emistry at the Polytechnic Institute of Ziirich

TRANSLATED AND REVISED

. BY

WILLIAM T. HALL, 8.B.,

Assistant Professor of Analytical Chemistry, Massachusetts Institute of Technology

Votume I

QUALITATIVE ANALYSIS

FOURTH ENGLISH AFTER THE EIGHTH GERMAN EDITION
TOTAL ISSUE, FIFTEEN THOUSAND

NEW YORK
JOHN WILEY & SONS, Ino.
Lonpon: CHAPMAN & HALL, Limtrep
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) 916

PREFACE TO THE FOURTH ENGLISH EDITION

THe German text upon which this book is based was written by
an American who has taught for many years at Zurich. The first
English edition appeared as an authorized translation by one who had
been teaching analytical chemistry for only three years at that time.
It was translated largely as a result of a chance remark of his assist-
ant, R. W. Balcom, who deplored the fact that the students could not
read German readily enough to make use of the German text as a
reference book. At that time, the translator was using, as he has
always used, the excellent book of A. A- Noyes as a laboratory
manual in Qualitative Analysis and the extremely useful text of

_H. P. Talbot for the preliminary work in Quantitative Analysis. The

results obtained by asking the students to purchase both volumes
of this book in addition have been exceedingly gratifying. Better
examination papers have resulted and there have been fewer unneces-

‘sary questions asked in the laboratory.

Recently Professor Noyes has greatly changed his text on Quali-
tative Analysis and this has unquestionably had considerable influence
upon the preparation of the fourth English edition of this. book. It
has been so thoroughly revised and so largely rewritten that it is no
longer fair to Professor Treadwell to publish the book as a literal
translation, although the writer remains in thorough sympathy with
Professor Treadwell’s views and does not wish, in any way, to dis-
claim the great benefit and inspiration he has derived from close study
of the original text. He must, however, express his -obligation to
other texts, particularly to those of Noyes, Stieglitz, Béttger, and
Ostwald, from which many of the ideas introduced into this text have
been copied. The general plan of the book has been kept the same,
but greater stress has been laid upon the theoretical side of the subject,
particularly with regard to the applications of the mass action prin-
ciple, the ionization theory, and the theory of oxidation and reduction.

The translator wishes to acknowledge his indebtedness to Mr.
Donald Belcher, who has read all the proofs of this edition and
offered many valuable suggestions.

Wiuu1amM T. Hatt.

MassacHusetts InstiTuTE or TECHNOLOGY,
February, 1916.

iii

frre f)

PREFACE TO THE FIRST EDITION.

Havine been repeatedly requested by former pupils to publish
the lectures on Analytical Chemistry given by me at this Institute
since 1882, and not having time then to do it myself, I permitted
the ‘“‘ Verein der Polytechniker ” in 1885 to print in manuscript form
the notes of one of my students.

This output met with such a friendly reception that in 1888 a
second edition became necessary. Subsequently I decided to revise
the material thoroughly and publish it in book form; this text-book
of Analytical Chemistry represents, therefore, a somewhat amplified
repetition of my lectures.

The book is intended not only for laboratory use, but also for

self-study. With each element the mineralogical occurrence, crys-
- talline form, and isomorphous relations are briefly mentioned. Then,
after explaining the reactions, the methods of separation are given
in the form of tables; because, contrary to the views of many, I have
in this way obtained the best results in teaching. These tables are
summarized charts by which the student can quickly find .his
bearings. :

Much weight is placed upon the determination of the sensitive-
ness of the single reactions, as explained on page 75, because the
beginner becomes in this way at once familiar with the solubility
of the most important salts, and also with simple stoichiometrical
calculations. The approximate solubility of potassium chloroplatin-
ate, for example, is found from the following determination of the sen-
sitiveness of the reaction by which it is formed:

If 100 cc. of the solution contain 0.156 gm. potassium, one finds

that the formation of the chloroplatinate, at ordinary temperatures,

only takes place on addition of a little alcohol; but on increasing

slightly the amount of potassium in the solution, it takes place imme-

diately. We can, therefore, assume that the solution, which con-
Vv

vi PREFACE TO THE FIRST EDITION

tains 0.156 em. of potassium per 100 cc. water, is saturated with chloro-
platinate; hence the amount of the latter may be calculated:

Ke : KePtCle =0.156 : x;
78.3 : 485.8 =0.156 : a;
z=0.97.

The result shows that 100 cc. of water, at ordinary temperatures,
dissolve 0.97 gm. of K2PtCle, while accurate determinations at 20° C.
have given the value 1.12. The difference, about 12 per cent, is.
explained by the facts that we did not work at exactly 20° C., nor
with absolutely pure water; the solution also contains an excess of
chloroplatinic acid, whereby the solubility of the potassium chloro-
platinate is diminished; evidently the values obtained in this way
permit a very good comparison of the solubilities of the different
salts. From the sensitiveness of the reaction between a potassium
salt and tartaric acid, the solubility of the potassium acid tartrate
may be found to be 0.38; so that the solubility of the potassium
chloroplatinate is to that of the potassium acid tartrate as 0.97 : 0.38;
the potassium tartrate is about three times as insoluble as the chloro--
platinate, etc. }

The size of the book does not permit going into the microchemi-
cal detection of the different elements. We have, however, in the
excellent work of H. Behrens, ‘ Anleitung zur mikrochemischen
Analyse,’’ a reference book of the highest rank.

In publishing this, the first volume of the work, I beg of my col-
leagues and fellow chemists to kindly inform me of any errors or —
omissions.

F. P. TREADWELL.

Zuricu, April 29, 1899,

INTERNATIONAL ATomMIc WEIGHTS, 1916

Symbol | eight Symbol | Atomic
Aluminium. ...... Al 27.1 Molybdenum.....| Mo 96.0
Antimony........ Sb 120.2 Neodymium...... Nd 144.3
PII 6 OS A 30.88. ll Neots cece Ne 20.2
VAPORS, 82. ac. S52 As 74.96 || Nickel........... Ni 58.68
DAMM Ss esse vs Ba 137.37 || Nitrogen......... N 14.01
Pema. yt... Bi 208 .0 Osnaium. 665. 22s Os 190.9 —
vt Satake Apel B 11.0 Ceyeon Obs O 16.00
Bromine.......... Br 79.92 || Palladium........ Pd 106.7
Cadmium. ....... Cd 112.40 || Phosphorus...... P 31.04
Caesium.:........ Cs 132.81 Platinum........ Pt 195.2
Calotum: .. ... 5.22; Ca 40.07 || Potassium....... K 39.10
CeROORy oes, ee C 12.00 || Praseodymium. . Pr 140.9
SOO a Se ee Ce 140.25 |} Radium: 2.0. 5... Ra 226.0
Chlorine... ...2..: Cl 35.46 || Rhodium........ Rh 102.9
Chromium........ Cr 52.0 Rubidium........ Rb 85.45
IOAN ais. toon oa Co 58.97 Ruthenium....... Ru 101.7
Columbium....... Cb 93.5 Samarium........ Sa 150.4
ee A ae lines Cu 63.57 || Scandium........ Se 44.1
Dysprosium...... Dy 162.5 Selenium......... Se 79.2
LS 0) 0 a Er 167.7 SINGOR ess. Pa sees Si 28.3
Europium........ Eu 152.0 SUG aaah ate Sie Ag 107.88
Fluorine.......... F 19.0 Sodnanhis o> 2s.40 8 Na 23.00
Gadolinium....... Gd 157.3 Strontium........ Sr 87.63
SL a ar Ga 69.9 Gaetars séo4% eens S 32.06
Germanium....... Ge 72.5 Tantalum........ Ta 181.5
Glucinum?!....... Gl 9.1 Tellurium........ Te 127.5
OE Rea Ne aA Pa Au 197.2 Terbiim. 34.5... Tb. 159.2
PIBMUIOS, i505 0. He 4.00 || Thallium........ Tl 204.0
Hydrogen........ H 1.008 || Thorium......... Th 232.4
pO a a In 114.8 Then 0 os es Tm 168.5
BOUMNOs: 05. 6 aie! 6 I 126.92 ‘Tihs Oe foes heen: Sn 118.7
a a ee Ir 193.1 ‘Pitaniine: S75 a8 Ti 48.1
LO aa ee eae Fe 55.84 || Tungsten........ W 184.0
MRerpion <=. <5. Kr 82.92 || Uranium......... U 238 .2
Lanthanum....... La 139.0 Vanadium....... V 51.0
OP Re eee arte Pb 207 : 20: Si Menos.: 2. eceeees Xe 130.2
PAGNUED f 22-3 olera'. Li 6.94 || Ytterbium.......
Lutecium......... Lu 175.0 (Neoytterbium) . Yb 173.5
Magnesium..... Mg 24.32 || Yttrium......... Yt 88.7
Manganese....... Mn $4.93. Zine. sia eee Zn 65.37
Mercury......... Hg 200.6 Zirconium....... Zr 90.6

1 Also called Beryllium, Be.

TABLE OF CONTENTS

PART I
GENERAL PRINCIPLES
PAGE
I NM Re SE Gol awa ic sie Sy Wadg ts Fas 4 lee tps v Frontispiece
International Atomic Weights, ROPERS gees es aE fos a 64 aS al RO, XV >
Qualitative and Quantitative Analysis............... 2c ces ce sce c eee ceeeees 1
RN ATU UE WU 20 sg ca da dS mro W an 3 ob '6 vod we, wtelbice'b ono Mca RD BTS 1
emGnEYy OF SuleCtrOly Ue DissOCintion. 2... ise cee bce e ene seas veneees 4
Mearention. OF Common MICCITOLYteS, . 2 ocak ids co veces csv cceaccscevdnkees 10
NINE AE CIES ROQEI dea g's cere iho eo die alba 6 a8 ove ere so wR a aa
Equilibrium between a Solid and a Liquid.................. 2. eee ee ee eee 11
Chemical Equilibrium and the Mass Action Law.................0.e0ee eee 13
Equilibrium between a Solid and Two Liquids..............0... 000-20 eee 17
Influence of Changes in Concentration upon Ionization...................--. 18
IEE IE REN sg ne Pe Ne Poros AE UES ehie yO ae aie cs Eee eee 19
NNR CME a eons WA Sg Ge a atohe seek me tak oS i dD cow og aie ob vb Og REE 24
IMEI NEC TENSES MONEE oe ot cx yar ss ee eA Ae, Ripe ee SEs pb Pane POP Rees 26
NURI HEIRS EME) PURMSIOTAONY 355)... stein Sh a eG sale Se Po bnle obo eos Fb ke eS 27
Electromotive Series and Oxidation Potentials.....................0020000- 37
ET Sab Th AMERTNONY LORE 55 65 css vcs ok ace Ss Wk alee Cae edible O40 40 eee eae 45
ERR Ne ot iy CN Scr enc oes SaaS OS cd 5 aie nb Ww ale oma ae 48
MSE EIENISE VOR 65 cea gia satiieg oie Ae wale sales so oben oe sei eee ead 52
Detection of Acids and Bases. ‘Theory of Indicators....................6.. 54
IER UNIR EE OMI rary c Micah tw Sing x sags Me 6 boa ee View ed bye Mae ohn ew 56
Putration and Waehing Precipitates... 0... 6c cee tent cnet ene e ewes 58
EOOUOIGAL DOMILIONS: nF ee cee ees BRON a URE Ge He Mas Ry Wa K cote Mie oone 58
ES CTE ROME WE PO, So De yk ge cle Vig rk dao Wp a mg a tiw Ste Mace 0 wie, bo edema 61
Mareen 15, Ci RO CAIS INUC GETOUIIG, 6 ka cs ek hele Ws bce o Slee sae es aeiseiobe 70
PN ML TSTTE, 078 SLE nt 55a bain kf hide a Gov mc Sewisc ks 8 ds wd lec bg alereece 71
Determination of Sensitiveness of Reactions... .........++ Os Wa Sx ca cee ee
PART II
REACTIONS OF THE METALS (CATIONS)
Group V (ALKALIES)
MINS ere tee gr Ooi al ertN paigd SEs fo a ws, ore e's AV Gee Wowie oe RTO 78
Re ra Nyy a Vis ee oie gh pic aretaiere Wisk de Mele ad go he LA CER MES 82
Ce Gl See Ae ALE OG of Fale cas #6 KET Kis WV as Sele pew 87
SRN NIN As ae oe eb fsb gia be 8 FINO side 6 Oia ie oles Leh PRES View eel n eee Ew Ew 93
Separation of the Metals of Group Vv eau hE Cab caine Oe db UCR ON ae se wep 96

x CONTENTS

Group IV (ALKauine Earrtus)

PAGE
General Characteristic Reactions. «.... 6 655 6.3 ek FAG pa eae Bebe vena coeee 101
ANNU 5 Sa 5 digs creo b ote abide trp ite piece ahate Ga bikes Be neite RIE oak a eae 102
Peramneiiem | ss ss Sak. Ss Sis de bo a a Hagin 3 me Wee be eee nae a eee 106
Ne a ne re on ai orem re SN er Cer SiN a ey 108
Separation of the Metals of Group IV............. cece cececcensvceveveees 110
Spectoscopic Analysis... .ccvslevvcrcvenseces ei iad» bad 4 Ora ele holon Se 115

Grovp III

PE MCUEASITIONIUL 5705 os so a'5. bc Sore 6 a wks aia Bd Ga URE a eh onie ea pee eke ala ae 125
MSPRPOMYMINATIN Sioa a6 6 cso ssc cen, Kp b OCRig bs oreee W pte Dlahaence Bibb he eeckeadece ok ene 132
OND ays icc 6 ¥a wb Ue we 6 Wa 0 bck wn e'm Fb URE Sta SPW A cscatener el ehal on Tegal ae 142
Gi k(t, ae Olen ERAS ERY eh eee SLR Pacem RES Oe 154
SR AUMTATURIYY 5.5 (a: 5 o-oo Gavin slb0ae’e Bin wih oes am emia sie M ace Sect a secalaile ib lm. 8 oye ee 157
RMI ATNN sod, 5 cos aia. 0 0:e preikie ace vi wd alnatg ie plata lalate a Be eae ets tn 160
MIO. 65g 3k ois! obo pcwla.8's Coad we MAS Lat Ae pildd eeelininlecele accuse 172
MOE MING os 55-5. Oos ia 6; w' 5;0'p, Bin 'v 0 din Whabe mae wid 'e tle Soe Silence beib eciase, dia vs d1atasia a ee 178
DEES six: ib 'o sia baal oad Wows Wise te OOP ek CE ROS Be eh ea 184
Separation of Group III from Groups IV and V.............e cece ce eeees 188

Analysis of Group III in Absence of Phosphates, ...screeeererreverseereess 189

Grovp IT
BROROINT on. scion. naw sates si0o.9/08 2 GEL Tb Unie eaiy.s 6 a eho k Wee ka eae ne 195
MIR a Diao! at roe SR ar a hs cae Oca Ge oes (ba ets ace ae ae 205
MUORTAININ eo h.6se 8 ed 0 Waele ace 3 a ARS AMO OE dD iano ORE pa ee 210
oe, PEE ET Te CE LPP T STEEP ey ee 214
EEE CLEC E PET PPE Es ie PrN Res PL 221
EP ELO EOE TOL FORE TE EOL ET COPE e PS PET TEL CER Ss ee 223
MOY. 8S os ks Roe EUS Ra bk 08 CAE ae we a os Wh ka es Daa 242
Me Ba tio Vein sg sen'eio. 0 V8 ae bes AM Ee hb wid le aM RLR ede Galak Cee nd bie 0 ne 249
RAMUS IE Gia le s2't ae eb ab a Wu. b-acacedwi'e winie WAS wonlareerel Mmalon alta em aoe bila: Graeme ann 261
BMORETANTRNY 5-594. a 555s, Dea av s'0 048 ae 0 hoo eivice '0'8 6 Nala oe ENG ad COE oy ee 266
Separation of Gold from Platinum... 6.5.6.6 scabsesvectevceen.ve scenes emi 270

Aanivaie OF Group Ics) 9.4 00:9 bs 0b'ee doe 64-0 VEEN 008 SR 271

Group I
RUBS Fi ia ksaiw eno se ek niples o chao e's io G Reade o! slilaid ain atau 0 bikin © ana eaeeen a e 278
Mercury (Mercurous Compounds)... oi. swicsccccceccuvceneneevn eup Oaeee 201
BROS age keh ae eo civiee'p woin sao Olu ORs nd oN aN aan see lg ee 205
Aunbvaw of Group 1, + 5. «.> sav adh din he lale's:p’> Vast 4d wlesele 9 0h be ee eae ear 282
PART III

REACTIONS OF THE ACID CONSTITUENTS (ANION 8)

Division of the Acids into Groups... ........sccecececceccuceececsecesencs 284

CONTENTS xl

Group I !
PAGE
Hydrochloric Acid............. CW ibe SN ied gm aais ba LRA se bees patwardcw made 285
EN Si iine a's o\e Ka ad Had hoa ee lwo 0 PO Goctehan A ME Sig? Se PET OY Fo DL TN 2 291
EDINA COMB ki ai sr ala,g Bie gi wie ble dv.oe once WldAav n'y hoe via 0 wicks SERS tA SneRe 294
IRIE AUD rent sry een eo aN orig as Hoe gst are S hle a ay hoes emewee Ue 297
ERs vere CA aL aie Lis tek OE Ue hae bin in aed ee hig Deall epprn eeee 299
DOOD fare Ae Oy Mars pase AT Rew Ve DOSES kore vibe be aise oe pa 301
NG Say cee nots PERT ENN UNS AE a i chore nS ociealees * ner he aod 304
Detection of HCl, HBr, and HI in the presence of one another.............. 306
Detection of Halogen in the Presence of Cyanide.................0000eeeee 309
NMED RMR aol gb ara oa) Gene Oras aly alesis ccola oie ' Ad oes eMbined y tiee oh Oeio ER 310
nh ee SOMA ha aC SVK ow Cds Palo a als AE bias diy Haale eae e slbe 316
Detection of Nitrogen in Organic Substances.............0.000000 00 cece eens 316
Detection of HCN in Presence of HCl, HBr, HI,H,[Fe(CN),], Hs[Fe(CN).] and
MER a SE AN gee Coorg sa fe ig oe iva oi ka wip Aw 0 Sm nike Ai a o/b. wee 4 oianeeaie Bad 317
EN PM ea aso f, 5 Sec Ovi g-5-47e'e a Se el pete Wine Side a>, WAN ge vie aeRO 318 _
IR RAR IE elias tite et Sr Nt | Sd ge alana oak hai we ate etgien 321
RPM tira eee ER ee ee ik ota ah g SWAG Abe a pasha yO wa ea 323
Dection of HCNS in Presence of Ha'ogen and Cyanide....................4. 325
Detection of HCNS, H,[Fe(CN).] and H3[Fe(CN)g]..............0.0 2 eee eee 327
MTC VANIG ACI. fo. cds vce cwe es ey ew ren Sie WAN ark ea ge at Nis ¢ uh 6 ea 328
Grovp II
ENRON CG US eden ok St gius flay Yolk w avy. 9 viet eitus-a Poe Roo ai boa ia. a Rip aps ea 329
Hydrosulfuric Acid (Hydrogen Sulfide).............. ccc cece cece eee c eens 334
Re a. coos Ayo SSE RVG ive AUD FE SSRN vie eNO ks CSW pie rainray mse 339
SERIES IRE ela nel 2 DERE IN ers gh ty SO a Pe om ke ae A 341
IIR One NM LL a Dag PAI See ere gly samme abe ae elie 343
IRI FRG tees 5, os vee bie oe wy Pre Gees WO eae € 40S PAE 345
Grovp III
I ON sad ad onl slelenns Ga Peete SP OMAR bbe ed RGR ES VU ewoed owes 347
RTE RES RINGS RIDE SEAL Ae tas i gi Pay. 0 ag Ie De em aE ae 352
NS 2 Sag na gd arsig Gk GM 9. ST ae RW Lees bee eae rdees 356
I SS De esa ioral ee hin sak Pelee Fe Sees Foss dC eae fe be bide ben came 357
TRIE ee Ne te eh RoR ke 5 cole bs cab ace bes wach b sf Wray alate 360
TERRE ed sare isi a) ae gk as Be ar er waar Ee LL 362
Meer ait Ie Skates FREE S64 cas > pre Re 8 pint ha bus a epee TTS 366
UE ON ee Bek pera g 504-4 bain @ Wea le ovialk Mera e melee gale 369
SD SOCIETIES a ee A ee ee a SO aC eT I ae 372
TE ONE ES IS eh ek Sa ee Pa a Rare era aT 374
TO: ASME ov as 'sag pee AP ae ee ater aca wloy'a Da UK Wins ce tea Re ekeTorarals 375
Group IV :
EATERIES ANNES es, A cus fav alo Chips 0d hatdws weeds eee Re ee Chae Re EL Oe 377
MINIS coh, sun a'a Re aih cts SP aan Owed Cob US RENT aeN Cage Cem eM EK RES 382
PE PRMMIONSN TDN on caks DaP A nce ble Se vwloidly avo 4 6 SNA Sie, s es bie pe ORM e Wb U.N 6 oR ale 224

IU EE pee tos ola ain, ots bap bie 3 kee CO Fe ME ONCE RY COREE ee Be 229

xii CONTENTS

PAGE
Chrome Acids oo. oe een Cos oss se ROO Ee ea eke bie a Dee 135
Thinsulfurie Aeid.. . 3.553 on. ah Pawo ke te ee 387
Detection of Sulfurous and Thiosulfuric Acids in the Presence of Hydrogen
NEO 5s eo sien slg satan vle’s Chios wiv vie Wino cre aban Stas ae ean gree 390
Groupe V
DUAGRID ACIDS Sooo 5ga5 ony ols 08 Sn aeitig Ok ylaicio a ois tw blvd Seas eee 392
Detection of Nitric Acid in the Presence of Nitrous Acid.................-. 395
RBIOPIC ACID ook os hp eeaacde g Oa lo pve ls betes we dieses en ee 397
Detection of Hydrochloric, Nitric, and Chloric Acids in the Presence of One
PAMNGTNO Fo og oS 5 oy6. vn 0.8 pm WL plc Bn abe wk aoa teas ks 398
Preeti Se 5 eae es boo oo wm obo wos eae pine ne es 400
PCPRTIC CIO ACID 00's Sones coe bila bla Su sensi vib aS ehlaep ok bites ay baste Chet 401
Monopersulfuric Acid (Caro’s Acid). 2 .3..0).7.005 oa. eek oa een spe 403
Group VI
SUREE ONO ACHE i Boos said ce bk sg ARS Ep ee pees 404
BR VAIN UNONES PRI 5 oe ah dks nee PA oon wield aa Gly Pee ee eee a G0 aaa 406
FLVOTORUOSNCIG. ACI «5.5 554° 85 v4. tyes b ndeise hh wh oc owiehe OME Oe eS ee 411
Group VII
PRPS OUE save a'a's bpce'e PSR See 4, 6-bs05G OSG alba BURR Sle a aca Nie Saget ta an 413
‘Treatment of Insotuble-Bailicsted.. ...... co i a Wee sso baie 9s Oo a 417
MMR Ss ag on ln oo Sia 0 Ba OP) Blaise eee eetere ean wah a ek dnote apo eae aaa 420
PART IV
SYSTEMATIC ANALYSIS
Analysis of Solid and Non-metallic Substances... .........00 0c eeeeeeeeeues 423
Prelinimary. Examimation « . 660006556 5.6 a hae 0 5s ve Ba we OE ee 423
Bomucn of the Substamce. ¢ oo... os ee eo a ee eie es ee eauee ee ae ee 430
SIAC TIO 5.600 Sie ne ok vind bo cle ho VR A ee Oy bale ge 432
Methods of Getting Insoluble Substances into Solution. ................+005 436
Examination for the Metals:(Cations). .....05..60. oso. eo sk nee tie 440
Examination for the Negative Elements casein, we a Saee wie sells eeepc 446
Analysis of Alloys...... Sw wb mele SRUe RR Ro Le G5 Salm k cg eaUe Lae cra ee 452
Ried inh ee FOR aees aera 452
PART V

REACTIONS OF SOME OF THE RARER METALS

Group V (ALKALIES)

CMAN ea Peis os be oie ah wc bre phe bole Gal TE IONE VO eee oa viel ae 456
Bea AI iio is ds ec cs oo os So 00s Dee CER fore k «ae 457
Bt NE Ra nae CN Bean A oe Se dees) 7 eR Sg ee Pe 458

Detection of Lithium, Rubidium, and Cesium in the Presence of Sodium and
WNC oo ec en denis oy ARIA aes A REAR e ots aT ae See 459

le Ven ease a ee
- ‘ a aw a

CONTENTS xili
Group III

PAGE

NT teas oa Wis Cs v6 Sha) 8 0S vias’ Perec araiots ieiyze!<-k.s.a/ee Cra aie wate ora oanS 461
ME etc oP ape eg etl A PRR PWC age keto 2 462
IN Sear ec hts MEN re Sg ce Tia des x bi co 5 vn age a 6 edn olga cae ce 465
EE II ae Sa dials. 3-6 ale 0 baa hid Vivian 4.46 Av'o. co oboe Wan bree ors 467
ED errs ROE Loe iE thay ees tain Uo gw a 30 bs ole view unlobic buen 469
MN rach ce ae tp Eie Oc isrg @ aul ole/e Ugalns Sis & Lady vd vee de cat ehabwn voles 472
ME Nn CG vig ho 2 eam ha Oey ods be bh LSG SL ede hab eee 473
II UES TAMUNEINAILO © 500° 5"5 Gib a < p.o 0:6 0b arco ofc calc Slee cdvicecscudncsvarensibac. 478
Ih ei tray Pe ee tals Soe ep ain sb a ee ieee Gee Oa Cow eae 480
Ne Nowe nedee naeeeucaen wae
Separation of Tantalum from Niobium..................seccceccscceseces 484

Group II

ICE DED Riest ye, et ON Ne ey ee ey a de hin eo h,0 0.4 oe aed vas CaaS 485
Re ae I oe Ee SY ad ae yiare. da ca wigte Neral he Waa Oba wa 487
MIE re en ens vc PSA Sh to 4x PGA PoNin wie es 6% ha Reine 491
RT ett fa a gre s-\\clan tna a wlarseh § acece Xb 5, bale ce bre tym es ween ee VES veer * 494
EE Rin ree et MEET Pe ie 5 ded win il kw acm w sle's GE Oop a 496
I RR SERS a TROL PPAR Aa ae” Ry Se pilpeo ry Ca RE aE ORE ROD? 499
CM toa he ee 28 Rb ca oaxe. aie. Pare chee FECL , BACT eS es, 504
REI De cr Te etre Le ol BLS rm ie me arg ae one Sy SM Ee wig ne Oe Bieid Chee 507
IE CO Bien SSL, Meant means Sure Gre try sl ajeia We Wie 5.4 o/s’ chawig biog areca Ons 509
NEN Le as et Cue as Beg Ane Cae e 6 VERE Oe bee 8 Uwe de eke 510
MER EAT vier STS ood IER ha ly lke MAA Y inky A eb bow a! 0 oo ee eS 512

penaration of the Platinum Metals... .cvrccsicccsneceeeseveevereqvensecns 516

QUALITATIVE ANALYSIS

PART I. GENERAL PRINCIPLES

By Chemical Analysis is understood all those operations which
are performed in order to determine the constituents of a chemical
compound (or a mixture of chemical compounds). Chemical Analysis
is subdivided into Qualitative Analysis and Quantitative Analysis.

Qualitative Analysis treats. of the methods for determining the
nature of the constituents of a substance, while Quantitative Analysis
treats of the methods for determining in what proportion the con-
stituents are present in any compound or mixture of compounds.

In order to recognize a substance we change it, usually with the
help of another substance of known nature, into a new compound
which possesses distinctive properties. This transformation we
call a chemical reaction; and the substance by means of which the
reaction is brought about, the reagent.

We distinguish between reactions in the wet way and reactions
in the dry way.

I. REACTIONS IN THE WET WAY

For the purpose of qualitative analysis only such reactions are
applicable as are easily perceptible to our senses. A reaccion is
known to take place—(a) by the formation of a precipitate; (6) by
a change of color; (c) by the evolution of a gas. In other words,
the sense of sight is used chiefly in qualitative analysis and most of
the reactions employed are visual ones. The sense of smell also
aids in identifying many substances. Thus the vapors of hydrogen
sulfide, hydrogen cyanide, bromine, carbon disulfide and a great
many other substances have very characteristic odors. Some of these -
vapors are poisonous, so that in trying the odor it is best to gently
waft a little of the vapor, by a motion of the hand over the substance
to be tested, in such a way that the vapor reaches the nostrils greatly

2 GENERAL PRINCIPLES

diluted with air. The sense of taste is sometimes useful, but is rarely
employed on account of the danger of poisonous effects. The sense
of touch sometimes furnishes a little aid; thus graphite has a peculiar,
greasy feeling, and paralysis of the tongue or eyelid is temporarily
imparted by the alkaloid cocaine and certain allied substances.

When an aqueous solution of barium chloride is mixed with dilute
sulfuric acid, a white crystalline precipitate of barium sulfate forms:

BaCle+ H2SO4 = 2HCl1+4 BaSO..

A precipitate of identically the same chemical composition can be
formed from any other soluble barium salt or by using a solution of
any soluble sulfate instead of sulfuric acid.

The addition of a little silver nitrate to an aqueous solution of
barium chloride causes the formation of a white, curdy precipitate
of silver chloride which darkens on exposure to light:

BaClo+2AgNO3=Ba(NOs)o+2AgCl.

The same precipitate is formed when hydrochloric acid or any other
chloride is used instead of the barium. chloride and when any other
soluble silver salt is used instead of silver nitrate.

In the same way there are certain properties which are shown by
aqueous solutions of all acids. Blue litmus is turned red, carbonates are
decomposed with effervescence, and metals are dissolved. These so-
called acid properties are due to the hydrogen of acids, which behaves
in an essentially different manner than the hydrogen of other com-
pounds.

Bases also show certain characteristic reactions which can be
traced to the hydroxyl, OH, that they contain. An aqueous solution
of a base turns red litmus blue and reacts with the hydrogen of an acid
to form water. :

The aqueous solutions of acids, bases and salts, therefore, show
reactions which are characteristic not so much of the dissolved sub-_
stance as a whole as of its constituents. This is a very important
point. It enables us to test for the constituents of a solution more
or less independently of what. other constituents may be present. We
can test for barium in just the same way whether it is present as
chloride or as nitrate, and we can test for chlorine by the same reagent
no matter whether the chlorine was originally present as hydrochloric
acid or as some other chloride. This is quite remarkable, because the
chemical properties of a compound are usually quite different from the
sum of the properties of its constituents. The properties of the chemi-

‘

REACTIONS IN THE WET WAY 3

cal compound water show little similarity to the properties of either
hydrogen or oxygen gas. The properties of sodium iodide are quite
different from those of metallic sodium and of free iodine, and those of
potassium chlorate are quite distinct from the properties of the potas-
sium, chlorine and oxygen which it contains. Aqueous solutions of
acids, bases and salts, however, actually do show additive properties,
i.e., sodium chloride in solution shows properties which the sodium
of any other sodium salt will show, plus other properties which any
other chloride will show. This suggests the hypothesis that the
aqueous solution of an acid must contain the acid hydrogen, to some
extent at least, in the same condition as in the aqueous solution of
any other acid; that an aqueous solution of a base must contain a
part at least of its hydroxyl in the same condition as the aqueous
solutions of any other base; and that the metals and non-metals of salts
must be present in very much the same condition irrespective of the
nature of the original salt. This would mean that when the acid,
the base, or the salt is dissolved in water, it is decomposed to some
extent into smaller units.

Not only the chemical behavior of aqueous solutions of acids,
bases and salts indicates that the constituents are present in a con-
dition such that they may react independently, but also the physical
behavior of the solutions. The boiling-point of a solution of sugar
in water is higher and its freezing-point lower than that of pure water.
It has been found that the rise in boiling-point and lowering of the
freezing-point is proportional to the number of molecules of dissolved
substance present. This rule holds so exactly for solutions of organic
substances dissolved in organic solvents that it serves for the deter-
mination of molecular weights. When, however, it is attempted to
determine the molecular weight of an acid, a base, or a salt, by deter-
mining the boiling-point or the freezing-point of its aqueous solution,
it is found that the molecular weight thus found is always too small.
In other words, a study of the boiling-point or freezing-point of acids
bases and salts indicates that the original molecules of the acid, base
or salt have been more or less split up into units smaller than the orig-
inal molecule. :

Finally the electrical behavior corroborates this view. It is well
known that substances behave differently toward the electric cur-
rent; some are conductors of it and others are non-conductors.
Metals are good conductors and sulfur is a non-conductor. Again,
the conductors are divided into two classes. Metals belong to the
first class and conduct electricity without experiencing any change
except that they become warmer. Conductors of the second class

4 GENERAL PRINCIPLES

are chiefly aqueous solutions of acids, bases and salts. Simultaneous
with the conduction of the current they undergo a chemical change, |
and decomposition products are obtained at each electrode.

Theory of Electrolytic Dissociation

If we insert between the poles of an electric battery a piece of
rock salt or some pure distilled water, there will be no electric
current in the circuit; a piece of fine platinum wire placed in the
circuit will not be made to glow. The solid rock salt and the dis-
tilled water are non-conductors of electricity; they are non-electro-
lytes: If, however, we dissolve rock salt in distilled water, and then
insert the solution between the poles of the electric battery, the
platinum wire will be brought to a bright glow, showing that the
salt solution is a good conductor of electricity— it is an electrolyte. It
is thereby proved that by dissolving the non-conducting rock salt
in non-conducting water an essential change of the former has taken
place. We can make the same observation with all acids, bases
and salts. In an anhydrous state they are non-electrolytes, while
in aqueous solution,* on the other hand, they are electrolytes. These
phenomena are readily explained by the theory of electrolytic dis-
sociation proposed by Arrhenius{ in 1887. According to this
theory, all electrolytes are partially decomposed in aqueous solution —
into electrically charged atoms or atom-groups called tons; and the
extent of this dissociation increases with dilution, until with very
great dilution it is practically complete. For every degree of dilution
there exists a certain state of equilibrium between the ions and undis-
sociated molecules. .

When the non-electrolyte rock salt is dissolved in water, it breaks
up, according to the equation

NaCl @ Na*+Cl-

into positively charged sodium ions and negatively charged chloride
ions. {

All salts, acids, and bases behave like rock salt. Thus sodium
sulfate decomposes according to the equation

NaeSO4u @ Nat+Nat+S0O¢-,

* They are also electrolytes in the fused state.

t Z. phys. Chem., 1, 631.

t Many chemists prefer to designate the positive ions by small dots and
the anions by small inclined dashes. The above equation is then written:
NaCl = Na+Cl’.

THEORY OF ELECTROLYTIC DISSOCIATION 5

fy and sodium hydroxide. into
NaOH @= Nat+OH-.

By this theory of electrolytic dissociation the phenomena of elec-
trolysis may be explained. very simply: If we insert the two poles of a
scurce of electricity into an electrolyte, one of the poles, the anode,
is charged with positive electricity, and the other, the cathode, with
negative electricity. The electro-positive anode repels the electro-
positive ions (cations) and attracts the electronegative ions (anions); _
and the latter, as soon as they come in contact with the anode, give
up their negative electricity, become neutral and separate out.* The
same thing happens at the cathode, where the electro-positive ions
(cations) are discharged. The amounts of electricity which are neu-
tralized at the electrodes are always renewed by the source of the
current, so that the process is continuous.

The electric charge on one atomic weight in grams of a univalent
ion is 96,500 coulombs ;t on an atomic weight in grams of a biva-
lent ion the charge is twice as much, and on a trivalent ion three
times as much. To deposit one atomic weight in grams of silver at
the cathode, therefore, it is necessary for 96,500 coulombs of elec-
tricity to pass through the solution, and there will be a simultaneous
discharge of an equivalent weight of anion at the anode. One
.coulomb is the quantity of electricity which is represented by the
flow of 1 ampere for one second, 96,500 coulombs, therefore, represent
96,500 ampere seconds or 26.8 ampere hours.

The transport of electricity in aqueous solutions takes place only
by means of the ions; the undissociated molecules take no part in the
process. The concentration of the ions and the conductivity of the
solution are quantities which are proportional to one another. It
is possible, therefore, to determine the extent to which a solution is
dissociated into its ions by measuring the electrical conductivity of
the solution.

The laws governing electrolysis were well understood by Michael
Faraday in 1834, and he gave the name of ions to those parts of the
solution which migrate toward the electrodes (cf. p.-10). The
positive electrode is called the anode and the negatively charged ion
which is attracted toward it is called the anion; the negative electrode
is called the cathode and the positively charged ion which is attracted

* In many cases the discharged substance at once reacts with the water, forming
ions again with evolution of either hydrogen or oxygen gas; these gases are, there-
fore, secondary products of electrolytic action.

{ This quantity of electricity is called one Faraday.

6 . GENERAL PRINCIPLES

toward it is called the cation. In Faraday’s time it was thought
that the first action of the electric current was to decompose the
molecules of the substance into the ions. About 1885 Arrhenius
made the simple observation that all those solutions in which the
dissolved substances have abnormally low molecular weights, as deter-
mined by boiling-point elevation, by freezing-point lower'’ng, or by
some similar method, are solutions which permit the passage of the
electric current—they are electrolytes; while solutions which give nor-
_ mal results in the determination of the molecular weight of the dis-
solved substances are non-electrolytes. Since 1885, therefore, it has been
believed by most chemists that electrolytic dissociation, or ionization,
takes place when an acid or a base or a salt dissolves in water,

This accounts for the fact that aqueous solutions of all silver salts
show similar reactions. They all contain the silver cation, and the
silver cation is different from ordinary metallic silver chiefly on account
of the fact that it bears a large electric charge. Most of the reac-
tions of qualitative analysis are carried out in aqueous solutions with
electrolytes. Most of the separations employed and most of the tests
are by means of reactions which are characteristic of the ions. For
this reason, a proper understanding of the theory of electrolytic dis-
sociation is necessary in the study of qualitative analysis.

Let us interpret the action of a dilute solution of an acid upon a
dilute solution of a base. In a dilute solution of hydrochloric acid,
for example, the hydrogen chloride is almost completely ionized, and in
a dilute solution of sodium hydroxide the base is also almost com-
pletely ionized. Hydrochloric acid and sodium hydroxide react
together to form water, which is itself but very slightly ionized. The
reaction between the dilute solutions of ila ys acid and sodium
hydroxide may be written:

H*+Cl~+Nat+OH~ =H20+Nat+Cl-.

By subtracting the ions which appear on each side of the equality

sign, the equation becomes:
H*t++OH~ =H.0O.

According to this, the neutralization of a dilute solution of an acid by
a base is merely the reaction of hydrogen ions with hydroxyl] ions to
form undissociated water. This is known to be true, because if the
reaction takes place with 1 gm. of hydrogen and 17 gms. of hydroxy] the
heat evolved is 13,700 calories. This same amount of heat is evolved
when an equivalent amount of a dilute solution of hydrochloric acid
is neutralized by a dilute solution of potassium hydroxide, or when
the hydrochloric acid is replaced by another acid such as nitric acid;

THEORY OF ELECTROLYTIC DISSOCIATION t

it represents merely the heat of formation of a molecular weight in
grams (one mole) of water from hydrogen ions and hydroxy] ions.

Similarly it can be shown that when an acid acts on a metal with
the formation of a salt and liberation of hydrogen gas, the quantity
of heat which is developed depends only on the nature of the metal
and is independent of the acid. The anion of the acid really does
not take part in the reaction at all.

The main assumptions of the Arrhenius theory of electrolytic dis-
sociation are as follows:. When an acid, a base or a salt dissolves in
water its molecules are immediately dissociated to some extent into
smaller fragments called ions. These ions are charged with electricity
and the sum of the positive charges residing on the cations is exactly
equal to the sum of the negative charges residing upon the anions
and the whole solution is electrically neutral. The dissociation is a
reversible reaction and all electrolytes may be considered to be com-
pletely ionized at infinite dilution. Except for the dependence
resulting from the electrical charges and the consequent attractions
and repulsions between ions, the ions may be regarded as independent
constituents with individual and specific chemical and physical proper-
‘ties. When a substance dissolves in water and is only partly dis-
sociated, then when the ions are removed, either by electrolysis or
as a result of*chemical reaction, the substance will at once dissociate
again to form new ions.

While it is true that nearly all acids, bases and salts are ionogens,
yet the extent to which the ionization takes place when the substance
is dissolved in water varies greatly. Thus a molecular weight in
grams of hydrochloric acid dissolved in 10 liters of water yields about
seventy times as many hydrogen ions as an equivalent quantity of
acetic acid; a similar comparison can be made with regard to sodium
hydroxide solution and ammonium hydroxide. In round numbers,
hydrochloric acid is about seventy times as strong an acid as acetic
acid and sodium hydroxide or potassium hydroxide is nearly seventy
times as strong a base as ammonium hydroxide.

On the other hand, a molecular weight of acetic acid will neutralize
the same weight of sodium hydroxide that a molecular weight of hydro-
chloric acid does, and a molecular weight of ammonium hydroxide will
neutralize the same weight of acid that a molecular weight of sodium
hydroxide does. In a solution of sodium hydroxide and of hydro-
chloric acid of the above concentration the original molecules are
about 90 per cent ionized, and when the acid and alkali are mixed
the principal change is the union of hydrogen ions and hydroxyl ions
to form water. When acetic acid of the same concentration is used,

8 GENERAL PRINCIPLES

there is present at the start only 1.3 per cent of all the hydrogen in the
form of ions. These ions will at once react with hydroxyl ions
to form water, but there is always a tendency for the acetic acid to
dissociate, and when the ions disappear as fast as they are formed
the ionization continues and soon all of the molecules of acetic acid
will have dissociated. In the neutralization of acetic acid with
sodium hydroxide, the final heat effect will not be simply that of the
union of hydrogen ions with hydroxyl ions, but will also involve the
energy required to cause the acetic acid to dissociate. When a sub-
stance ionizes as soon as it dissolves, the heat effect of ionization cannot
be distinguished easily from the heat of solution. Just as some sub-
stances dissolve with absorption of heat and some with evolution of
heat, so it is found that the ionization process may likewise be asso-
ciated with either an absorption or evolution of heat.

It is interesting to note, and this is a matter of considerable
importance, that the salts of weak acids and of weak bases are usually
ionized nearly as much as the salts of strong acids or of strong bases.

When a bivalent acid dissolves in water, the two hydrogen atoms
do not dissociate to an equal extent. The ionization takes place in
two stages. Thus with sulfuric acid the first stage takes place in the
sense of the equation: ' |

H2804 @ H*t+HS0O,-.

The fact that the reaction does not necessarily take place completely
is indicated by using the double arrow sign instead of the equality
sign. When the above reaction stops there is a state of equilibrium
between the three substances H2SO4, H* and HSO4~. The HSO,-
undergoes a secondary dissociation as follows:

HSOs. = Ht+S0.-.

The extent to which these reactions takes place depends entirely
‘upon the dilution. If half a molecular weight in grams of sulfuric
acid is dissolved in 10 liters of water, the primary dissociation will
take place to about 90 per cent of the entire quantity of acid present
and the secondary dissociation to less than 50 per cent. If the solution
is extremely dilute, both reactions will take place almost completely.

In the case of carbonic acid, the primary stage

H2CO3 @ Ht+HCO;-

ordinarily takes place only to a fraction of 1 per cent and the second
stage
HCO3 @Ht+CO3=

THEORY OF ELECTROLYTIC DISSOCIATION 9

to an inappreciable extent (cf. p. 10). With hydrogen sulfide the
relations are similar.

On the other hand, the salts of these weak acids will dissociate
almost completely as follows:

NazCO3 = Nat+Nat+C03*
NazS = Nat+Nat-+S-*.

The fact that the extent to which a substance ionizes is dependent
upon the concentration of the solution has already been indicated
(p. 4) and will be demonstrated mathematically a little. farther on
(p. 18). The concentration of a solution shows the quantity of dis-
solved substance present in a unit of volume, and the numerical
value representing this concentration depends entirely upon the
units in which the mass of the dissolved substance and the volume
of the solution are expressed. If w represents the mass of dis-
solved substance in grams and v is the volume of the solution

expressed in liters, then = is the concentration of the solution

in grams per liter. Concentrations are often expressed in these
units, but the weight is not the most convenient unit for ex-
pressing the mass, especially in matters of theoretical discussion.
A solution containing 1 gm. of dissolved substance A is rarely equiv-
alent to one containing the same weight of a substance B. It is much
more convenient to measure the mass of the dissolved substance in
terms of the number of molecules present. According to the sug-
gestion of Ostwald, the term mole.has been given to the molecular
weight of a substance in grams, and when the concentration of a solu-
tion is expressed in the number of moles present in a liter, the so-
called molal concentration is obtained. The objection still remains,
however, that one molecule of a substance A (e.g., hydrochloric acid)
is not always equivalent to one molecule of a substance B (e.g., sul-
furic acid). To overcome this difficulty, concentrations are often
expressed in gram equivalents per liter, using the univalent substances
as the standard. Thus a mole of hydrochloric acid is one equiva-
lent weight in grams and half a mole of sulfuric acid is one equiv-
alent weight in grams. A solution which contains one equivalent
weight in grams of dissolved substance is called a normal solution;
one containing two equivalents in a liter is a twice-normal solution;
and one containing half an equivalent is a half-normal solution. All
things considered, this is the best way of expressing concentrations.
The following table will be found useful in studying the disso-
ciation of electrolytes. It gives the approximate percentage ioniza-
tion of substances present in 0.1 N solution at 25°. In the case

10 GENERAL PRINCIPLES

of polybasic acids, the value opposite the formula of the acid shows
the fraction of the whole molecule which undergoes the primary
dissociation into one hydrogen ion, that opposite an ion with a
univalent charge shows the extent to which this ion undergoes a
secondary dissociation, and that opposite an ion with a bivalent charge
shows the extent to which it undergoes a tertiary decomposition, —
forming a third hydrogen ion from the original neutral molecule of the
acid,

- Ionization Values of Common Electrolytes*

In 0.1 normal solution

Per Cent
Salts of the type BTA (ez, KINO) Sea ae eee 84
Salts of the type B,+A= or B* t+A,7 (e.g., KoSO4 or BaCl)f............. 73
Salts of the type B;stA=, or Bt + +A;- (e.g., K3Fe(CN)s or AICI;)........ 65
Saltsiof the type BttA™ (e.@.. MgSO)... os ates sce so bs ap eee 65
Moyne Ss hy Ooh Ee SSeS ee 90
Ba(OH). ees EB syhtoly “eine Sie 6 fb. pb apauw Bele te Sues te wpe ah etihts Shae 5] Cue: e, lene, sees mea rae ee ae 80
18 | Sa a ee ae eee ote es ace MM ee Mica Ma vr a 1.3
HCl, HBr, HI, HSCN, HNO;, HCIO;, HCO, HSO,, HsCrO............ 90
H3PO,, H;AsQOu, H.2SOs, H.2C20,, HSO,— a) dhe betel gah oer oe ae BRL aE rae et Ea 20-45
MNT OSS i ee Ce Ct ae 7-9
HC.H;02, HC.0,, PSO ge ee nr eh ie Se ee eee 1-2
HS, H.COs, HPO, = ’ HCrO,~ 4L 5. Mellen fe fa hails Noe ASS RRA Snel SRS Ae hes) De te ica eens ee 0.1-0.2
PCa IA e, THON, BOO o.oo b sene as) <a Nica ae eee 0.002-0.008
Rs MES Fes te ee ee ise Reo cS a 00001-0002
SPR OTE BG BD alice Gh bee ihc Ha 8 6 obebay SARE wie AOR Fae RN oe ee se ere 0 .Q000002

Nomenclature of the Ions

As sieaeas mentioned (p. 5) Faraday in 1834 was the first to use
the words ion, cathode, cation, electrode, anode, etc. ‘These names are
all derived from Greek roots. Faraday’s idea was that the electricity
entered the solution at the positive pole and passed down to the
negative pole. The word zon is the Greek word for wanderer or traveler
spelt in Latin letters. Anode is from the Greek dvé (ana) up; 600s
(odos) a path. Cathode is from the Greek xara (kata) down; 600s
(odos) a path. The two electrodes are considered as the doors or
paths by which the current enters and passes out of the solution.

Two methods are in common use for designating the ions. Thus
the ions of hydrochloric acid are designated as H* and Cl” or as H’
and Cl’. Small plus and minus signs are used in this book rather than

*From A. A. Noyes, Qualitative Chemical Analysis.
t Exceptions: CdCl, ionizes to about 47 per cent, HgCl to about 0. 01 per
cent, and HgBr,, HgI., and Hg(CN)s less than HgCh.

y =

eve eh os

NOMENCLATURE OF THE IONS 11

the dots and dashes simply because this is the present practice in
the journals published by the American Chemical Society. Many
writers prefer to use the other system because it takes up less room;
in the case of the polyvalent ions the use of the plus and minus
signs is often very cumbersome.

Purely as a matter of convenience, an attempt has been made
to devise a system of rational nomenclature for theions. This method
has been adopted in a number of excellent text books but it is not

in common use. According to this system, the names of the cations

are obtained by adding the termination -ion to the stem of the name
of the corresponding metal, using the Latin name whenever possible.
When a substance forms several ions differing from one another only
in valence, the names of such ions are designated by Greek prefixes
indicating the number of charges residing on the ion. The names of
the anions are derived from the names of the salts. If the name of
the salt ends in -ate, these last three letters are replaced by the ending
-anion, except in the case of the carbonate ion, which is called carbanion.
The names of anions from salts ending in -ite are formed by replacing
these three letters with -osion. The anions from salts whose names
end in -ide are obtained by replacing these letters with the ending
-~idion. 'The hydrogen ion is called hydrion and the hydroxy! ion is
called hydroxidion. The following table illustrates the use of this
system which was proposed by Walker: 3

NAMES OF CERTAIN IONS

Symbol. Name. aces reg Symbol. Name. Anion of
Agt Argention Silver Cl- Chloridion Chlorides
Cat + Calcion Calcium ClO— | Hypochlor- Hypochlorites
Cut + Dicuprion Cupriec copper osion
Fet + Diferrion Ferrous iron ClO,;— | Chlorosion Chlorites
Fet++ | Triferrion Ferric iron ClO;—_ | Chloranion Chlorates
Ht Hydrion Hydrogen ClOg | Perchloranion | Perchlorates

(Acids) 33 Sulfidion Sulfides
Kt Kalion Potassium SO;= Sulfosion Sulfites
Nat Natrion Sodium SO Sulfanion Sulfates
NHgt Ammonion | Ammonium NO;— _ | Nitranion Nitrates
OH Hydroxidion | Hydroxides
(Bases)

Equilibrium between a Solid and a Liquid

Most of the reactions used in analytical chemistry involve either
the solution or the precipitation of some substance. It is important,

12 GENERAL PRINCIPLES

therefore, to consider briefly the. relations which exist between a solid
and its solution. :

Potassium nitrate on being brought into contact with water at once
begins to dissolve. The rate of solution is influenced somewhat by
the amount of surface exposed by the salt, a fine powder dissolving
more rapidly than a single large crystal. At first the substance dis-
solves quite rapidly, particularly if the liquid is kept stirred, but
gradually the speed slackens and finally a time comes when the water
at. a given temperature will dissolve no more of the salt. The solu-
tion is then said to be saturated with the salt and it makes no differ-
ence how much potassium nitrate is available in excess of the amount
required to form a saturated solution, the solution when once satu-
rated at any temperature will dissolve no more salt.

The quantity of salt required to form a saturated solution varies
with the temperature, more so with potassium nitrate than with
many other salts. At 0° the saturated solution contains only 1.3 moles
of potassium nitrate, whereas 2.7 moles dissolve at 20° and ” moles
dissolve at 100°.

If a solution of 5 moles potassium nitrate is prepared by dis-
solving the salt in hot water and the solution is then cooled to 20°,
we obtain what is called a supersaturated solution. A state of super-
saturation can be maintained for some time provided care is taken
not to disturb the solution in any way. If the supersaturated solu-
tion is agitated, or, better, if a tiny fragment of potassium nitrate is
thrown into it, crystallization starts and continues until finally the
solution only contains 2.7 moles of the salt, which is the quantity of
potassium nitrate required to form a saturated solution at 20°.

The solubility of a substance at any temperature is usually deter-
mined by two methods: first, by shaking up the salt with water until
a saturated solution is obtained; second, by forming a supersaturated
solution and allowing the excess of the salt to crystallize out. Usually
the values obtained by the former method are a little lower than the
values obtained by the latter method; a slightly undersaturated solu-
tion is obtained in one case and a slightly supersaturated one in the
other.

When a solution of potassium nitrate is brought into contact with
more of the salt, whether more of the salt will dissolve or not is
determined solely by the concentration of the solution. If it is
saturated already with potassium nitrate, no more of the salt will
dissolve; if unsaturated, more salt will dissolve to form a saturated
solution. The equilibrium between a liquid and a solid which is
soluble in it is determined solely by the concentration of the solution.

CHEMICAL EQUILIBRIUM AND THE MASS ACTION LAW 13

The absolute quantity of substance and the absolute quantity of solu-
tion have no effect upon the final equilibrium.

Chemists prefer to look upon a state of equilibrium as a condition
of dynamic equilibrium rather than as one of static equilibrium.
Instead of thinking of the saturated solution of potassium nitrate as
one which has no tendency to dissolve more potassium nitrate, it is
preferable to consider the solution as one in which the tendency to
precipitate potassium nitrate is exactly balanced by the tendency
to dissolve potassium nitrate. When the solution is undersaturated
and more salt is available, the tendency to dissolve is greater than the
tendency to precipitate and when the solution is supersaturated the
tendency to precipitate is greater than the tendency to dissolve.

The equilibrium principle is the same in the case of difficultly
‘soluble substances. It requires only 0.0015 gram (=0.01 millimole)
of silver chloride to form a saturated solution in water. If more
than this quantity of silver chloride is produced as a result of a
chemical reaction taking place in an aqueous solution, all the excess »
silver chloride will be precipitated. The solubility is so slight that
the precipitation is practically complete.

Chemical Equilibrium and the Mass Action Law

If hydrogen sulfide gas is passed into a solution containing zinc
chloride, a white precipitate of zinc sulfide is formed:

ZnClz+He2S = ZnS+2HCl.

If the precipitate of zinc sulfide is filtered off and treated with hydro-
chloric acid it will dissolve:

ZnS +2HCl1=ZnCl2+H28.

Similarly, the addition of ammonium carbonate to a solution of
calcium chloride in water causes the formation of a white precipitate
of calcium carbonate:

CaCl2+ (NH4)2C03 =CaCO3+2NH4Cl.

The precipitate can be dissolved, however, by boiling it with ammonium
chloride solution.

In each of the above cases, there are evidently two opposing
tendencies—the tendency of zine sulfide to precipitate and the tend-
ency of zinc sulfide to dissolve; the tendency of calcium carbonate
to precipitate and the tendency of calcium carbonate to dissolve. To
express the fact that the reaction may go in either direction it is cus-

14 GENERAL PRINCIPLES

tomary to writé the symbols separated by a double arrow instead of
by an equality sign (cf. p. 8):

ZnCle+H28 = ZnS+2HCl,
CaCle+(NH4)2CO3 @ CaCO3 +2NH.,Cl.

Such reactions are called reversible. It was once thought that
reversible reactions were of rare occurrence, but it is now customary
to consider all chemical reactions as reversible, although in many
cases and especially in most reactions used in analytical chemistry,
the reaction goes so completely in one direction that only a negligible —
quantity of one or more of the initial substances remains unchanged.
In general, when two substances A and B react with one another at a
constant temperature to form C and D, then, to some extent at least,
C and D react to form A and B, and equilibrium is reached when the
ratio of the product of the concentrations of A and B to the product
of the concentrations of C and D has a definite, constant value. This
value is characteristic of the equilibrium between the compounds
involved. | ;

In the above case, the reaction may be expressed as follows:

A+B 2C+4D,

in which A, B, C and D represent four different substances reacting
in the molecular proportions indicated by their symbols. The condi-
tions of final equilibrium is expressed by the mathematical equation:

[A] X[B]
[C]X[D]

in which [A], [|B], [C], and [D] represent the final concentrations of the
four reacting substances and k is some definite number called the
equilibrium constant. The value k varies with the temperature.

If more than one molecule of substance takes part in the reaction,
the conditions are somewhat more complicated. This is expressed by
the general equation

=k,

mA+nB = pC+¢qD
and the final equilibrium is expressed mathematically
[A] x[B]”
[C}? x[D]*
in which [A], [B], [C], and [D] represent, as before, the concentrations
when equilibrium is reached.

=k,

ee a ee a, eee eee See

CHEMICAL EQUILIBRIUM AND THE MASS ACTION LAW 15

This is the so-called law of mass action, which was discovered by

-Guldberg and Waage in 1867. It is to be noted that it is the concen-

trations, or masses present in a unit of volume, rather than the actual
masses of the substances, which find expression in this law.
This law apphes to a state of homogeneous equilibrium. A homo-

geneous system is one in which every part of it is like every other

part. A mixture of two solid substances is not homogeneous. A

solution, on the other hand, is homogeneous when it is thoroughly

mixed, as it is impossible to distinguish any difference between differ-
ent portions of the solution. Similarly a mixture of gases represents
a homogeneous system. Such homogeneous systems are called
phases. A mixture of a solid, a so‘ution and a gas represents three
phases; two solids, two phases; two immiscible liquids, two phases.

In the case of the reactions between zine chloride and hydrogen
sulfide and between calcium chloride and ammonium carbonate a
precipitate was formed in each case. The mass-action law applied
only to the zine sulfide and to the calcium carbonate that remained in
solution. The fact that these substances are only very slightly soluble
in water favors the progress of the reaction in the direction by which
these substances are formed. The mass-action law shovs that when
the concentration of any substance participating in a chemical reac-
tion is increased, this tends to increase the tendency for the reaction
to take place in the direction by which this substance is decomposed;
when any substance formed by means of a chemical reaction is removed,
this increases the tendency for the reaction to proceed in the direction
by which this substance is formed. The formation of a precipitate
or the escape of a gas, as fast as the substance is formed by means of
a chemical reaction, tends to make the reaction take place more com-
pletely. If the gas is all boiled off the reaction by which it is formed
will take place completely. Similarly if any precipitate were absolutely
insoluble in water, the reactions by which this substance is formed
would take place completely.

This law of mass action embodies one of the most important
principles utilized in analytical chemistry. It enables one to under-
stand why most of the reactions take place and to establish conditions
under which these reactions will occur to the best advantage. The
law has been verified by a great many quantitative as well as
qualitative experiments. It has been studied, for example, in
connection with the formation and decomposition of phosphorus
pentachloride.

When chlorine gas reacts with cold phosphorus trichloride, the solid
pentachloride is formed; but if this substance is heated, it breaks

16 GENERAL PRINCIPLES

down into its constituents. The reaction is reversible and may be
expressed as follows: |
PCls @ PCls3+Cle.

At any given temperature an equilibrium exists which can be expressed
mathematically, according to the mass-action law,
[PCl3] X[Cle] a
[PCls] :

in which [PCl3], [Cle] and [PCls5] represent the concentrations at the
time when equilibrium has been reached. -

If we desire to volatilize phosphorus pentachloride so that the
least possible dissociation will take place, the above equation shows
us how this may he brought about.

If either [PCls] or [Cle] be increased, then in order that the value
of the fraction

[PCls] [Cle]
[PCls]

remains constant, it is evident that the concentration [PCl5] must
become greater; or, in other words, the dissociation of the penta-
chloride becomes less and there will be practically no dissociation if
the pentachloride is volatilized in an atmosphere of phosphorus tri-
chloride or of chlorine. In this way, Wurz obtained for the density of
phosphorus pentachloride 6.80-—7.42, instead of the calculated value 7.2.

At Stassfurt the mineral carnallite (MgCle- KC]-6H2O) occurs and
was evidently formed by precipitation from solutions containing the
chlorides of magnesium and potassium. This double salt is less soluble
than pure magnesium chloride and more soluble than pure potassium
chloride. If carnallite is dissolved in water and the solution allowed
to evaporate until crystals are deposited, it will be found that the
crystals consist of potassium chloride. When the carnallite dissolves
in water, the double salt is decomposed, more or less completely accord-
ing to the dilution,

MgCle: KCl @ MgCle+ KCl,

and for every concentration the equation holds:

[MgClo] X[ KCl]

iMeClo aon: COnStane

If we wish to recrystallize the carnallite, the breaking down of the double
salt must be prevented as much as possible, and to do this it is merely
necessary to add an excess of MgCle. As a matter of fact, the mineral

EQUILIBRIUM BETWEEN A SOLID AND TWO LIQUIDS 17

is recrystallized at Stassfurt from a 23 per cent solution of magnesium

chloride.

This law of mass-action applies to all cases of chemical equilibrium
that takes place in a homogeneous phase, i.e., it can be applied to
all reactions which take place between gases and to all reactions that
take place in solution.

The law of mass action applies to the equilibrium between an ion-
ogen and its ions. In this connection, the law is of particular impor-
tance in the study of analytical chemistry.

Equilibrium between a Solid and Two Liquids

Although water, either pure or containing dissolved acid, is the
solvent most used in analytical chemistry, it often happens that a
substance is more soluble in some other liquid. Thus free iodine
is about 400 times as soluble in carbon disulfide as it is in water. -
When iodine is in contact with both carbon disulfide and water, and
these two liquids are only slightly soluble in one another, it will dissolve
chiefly in the carbon disulfide. Moreover, if an aqueous solution of
iodine is shaken with carbon disulfide, the latter, when it separates
out beneath the water, will contain nearly all of the iodine. A state of
equilibrium then exists between the solution of iodine in water and the
solution of iodine in carbon disulfide. Such an equilibrium is governed
by the so-called distribution law or law of partition. If C4 represents
the concentration of a substance in a solvent A and Cz is its con-
centration in a solvent B, equilibrium is reached when

C7
This is the mathematical expression of the distribution law. The con-
stant, k, is called the distribution coefficient. In this simple form,
it is important to note that the law holds only when each concentra-
tion is expressed in terms of the same molecular species. Thus if a
substance is dissociated to a large extent in one solvent and scarcely
at all another, the concentrations involved must be those of the un-
dissociated salt in each case. It is quite common to find that the
ions of a substance are much more soluble in water than in any other
solvent whereas for the undissociated substance the relations are
reversed. Iodine dissolves to a greater extent in a solution of potas-
sium iodide than it does in pure water, owing to the formation of KIg3.
In such a solution the following equilibrium exists:

KI+I, @ KIsz.

18 . GENERAL PRINCIPLES

If such a solution is shaken with carbon disulfide, the distribution
law holds only for the free iodine held in solution as such in each
solution.

For iodine in pure water and iodine in carbon disulfide, the dis-
tribution coefficient is z4y. Theoretically it is impossible to. remove
all the iodine from water by shaking with carbon disulfide, but if the
carbon disulfide is removed, with the aid of a separatory funnel, and the
aqueous solution is shaken with fresh carbon disulfide, it is evident
that the quantity of iodine remaining with the water is negligible, -
or can be made so by repeating the operation.

Sometimes in testing for the halides it is desirable to remove free
halogen from the aqueous solution; to accomplish this, the distribu-
tion principle is utilized. Ferric chloride is much more soluble in
ether and hydrochloric acid than it is in water and hydrochloric acid;
to detect the minor constituents of iron or steel, a large sample of the
- original material is taken and the ferric chloride removed by shaking
the hydrochloric acid solution with ether. Perchromic acid is more
soluble in ether than it is in water; by shaking the dilute aqueous
solution with a little ether, a concentrated solution in the latter is
obtained and the presence of the chromium shown by the beautiful
blue color.

Influence of Changes in Concentration upon the Ionization of
! Electrolytes

If we assume 1 gm. molecule of a weak electrolyte, such as ammo-
nium hydroxide, to be dissolved in »v liters of solution, the original
substance will be partly dissociated according to the equation

NH,OH @ NH4y*+0H-

into ammonium and hydroxyl ions. If a gram molecule of the base
is dissociated in the sense of the above equation, then the undisso-
ciated part will amount to 1—a.
The concentrations per liter are
| undissociated dissociated
NH40H @ NH4t+OH-

Se.

Sea ® v
and according to the mass-action law

we ae a eae hese aa ay é-
- 7 se iy Me

SOLUBILITY PRODUCT 19

The constant k is known as the tonzation- or affinity-constant;

_ it is independent of the dilution and is characteristic for every elec-

trolyte. The expression shows, however, that by increasing v, or
diluting, the fraction of the molecule dissociated (a) will be made
larger.

If to the solution of the base we add n additional ammonium ions,
by adding solid ammonium chloride, then if is considerably larger

than a the degree of dissociation of the base will be greatly diminished,

namely, from a to a1, a value which can be readily computed as follows:
In the solution there is present per liter

undissociated dissociated
NH4OH @ NH4t+0OH—
l-—a, ai+n a1
v v v’
therefore
ps (aitn)a1
: (1 —ay)v :
If k and n are known, a; can be computed:

or (n+vk) + V (n+vk)?+40k
za ; ;

In the case of 0.1 N ammonia solution the ammonium hydroxide
is dissociated only to an extent of 1.32 per cent, the dissociation con-
stant being 0.000018. If we add to 10 liters of this ammonia solution
2 gram-molecules of ammonium chloride (107.08 gms.), then since the
ammonium chloride is 93 per cent dissociated at this dilution, we are
adding 2X0.93=1.86 NHa4 ions. If this value be inserted in the
above equation, ai becomes 0.00009; in other words, the dissociation
of the ammonium hydroxide is diminished by the addition of the am-
monium chloride from 1.32 to 0.009 per cent. The solution now
contains so few hydroxyl ions that it will not cause precipitation in
solutions of magnesium salts (cf. pp. 46, 94).

Similarly the dissociation of weak acids is lessened by the addi-
tion of their salts. In the case of the stronger acids and bases, the
effect of adding a neutral salt to the solution is not so remarkable,
because the stronger acids and bases are dissociated to about the
same extent as their salts.

Solubility Product

Silver chloride is slightly soluble in water; 0.00001 mole (1.5
mgs.) of the solid dissolves in 1 liter of the solvent. When water

20 . GENERAL PRINCIPLES

is placed in contact with an excess of silver chloride, a state of
equilibrium is soon reached between the solid and the solution. If
more than this quantity of dissolved substance is present at any time,
the solution is supersaturated and tends to precipitate silver chloride;
if less, then more silver chloride will be dissolved. When equilibrium
is reached the tendency of the salt to precipitate is equal to the tend-
ency of the salt to dissolve.

The dissolved silver chloride also exists in a state of equilibrium
between the ionogen and the ions. This equilibrium apparently
takes place almost instantly, whereas the equilibrium between the
solid and the solution is established more slowly. For all concen-
trations of such a slightly soluble substance as silver chloride, it is fair
to assume that the mass action law holds rigidly. Applied to the
reaction

AgCl @ Agt+Cl-

and denoting the concentration of non-ionized silver chloride as [AgCl],
of silver ions as [Ag], and of chlorine ions as [Cl], the law is expressed
as follows:

[Ag] X[Cl]

[AgCl] =f

In this equation k has a definite value, called the ionization constant,
which varies with the temperature but has otherwise a definite value
for every substance. When the solution is saturated with silver
chloride the value of both numerator and denominator has reached
the saturation value. If the value of [AgCl] ismade greater than
corresponds to this saturation value, the solution is supersaturated
and is no longer in equilibrium with undissolved solid. The value ©
[AgCl] could be used for expressing the solubility of the substance,
but it is often more convenient to use the ion concentration product
[Ag] X[CI] which is called the solubility product. In general, if the sub-
stance A,,B, ionizes into mA and nB ions, the solubility product,
Sy, is found by the following equation:

= [A]”x [B}” .< k[|A,B,l,

in which k is the ionization constant and the concentrations are those
of a saturated solution.

Experience has shown that the conditions are somewhat more
complicated in concentrated solutions such as are obtained with the
very soluble substances. In future discussion, therefore, the solu-
bility will be expressed, as a rule, in terms of the solubility product

a ee es ae o = ee

SOLUBILITY PRODUCT | 21

; only when the substance does not dissolve to a greater extent than

0.01 mole per liter. The table on page 10 shows that binary salts

of the type represented by AgCl are ionized to about 84 per cent
in 0.1 N solution, and we have seen on page 19 that the ioniza-
tion increases as the solution is diluted. In the case of such a dilute
solution as that of silver chloride (0.00001 normal) the ionization
is nearly 100 per cent. It is therefore logical, in such cases, to express
the solubility in terms of the ions, whereas in the case of the very solu-
ble substances it is better to measure the solubility in terms of the
mass of dissolved substance. | .

The numerical values of the ionization constant and of the solubility
product depend upon the units used in measuring the mass of the
dissolved substance or ions and in measuring the volume of the solu-
tion. It is customary, in this case, to express the concentration in
moles per liter (cf. p. 9). In the following table the solubility of the
substance is expressed in three ways: in grams of dissolved substance
per liter, in moles per liter and, finally, in terms of the solubility prod-
uct, using moles per liter. The solubility of most of the substances
given in the table is so slight that the quantity dissolved is negligible
for most purposes. Whenever the word insoluble is used in this book
it is with the understood limitation that no substance is absolutely
insoluble in water.

SOLUBILITIES AND SOLUBILITY PRODUCTS AT ROOM
TEMPERATURE

Substance. eames, a Fc pompscaty SoA Solubility Product.

JS eee 1.1X10-4 | 5.9X10-7 | [Ag]x[Br]=3.5x10-"
Ag.(CN)s...... 4.3X10-5 1.6X10-7 | [Ag]x[Ag(CN).|=2.6X10-"4
AgCNS....... 1.4x10-4 8.4X10-7 [Ag] X[CNS] =7.1x10-#
an 1.5X10-3 1.1X10-5 | [Ag]x[Cl]=1.2 10-1
AgeCrO........ 2.5X10-? 7.5X10-5 [Ag]? [CrO,] =1.7x10-™"
AgeCrO7...... 8.3xX10-? 1910-4 [Ag]? X[Cr2O7] =2.7x10-"™
7 ee 2.1X10-2 | 9.0x10-5 | [Ag]x[OH]=1.9X10-8
Oo SEs 3.010-8 1.3X10-8 | [Ag]x[I]=1.7x10-16
ST: ae 4.4X10-2 1.5X10-* | [Ag]x[I03]=2.310-8
MPO we... 6.5xX10-3 | 1.6X10-5 | [Ag]*x[PO.J=1.8X10-18
MIMO. 5 8.0 2.6X10-2. | [Ag]?<[SO.]=7.0X10-5
ES Se eae 1.3X10-? 4.3X10-5 [Ba] X[CO3] =1.9 10-9
BaCrQy....... 3.8X10-* | 1.5xX10-5 | [Ba]x[CrOJ=2.3x10-%
DOE. Cis es 5. 2.5X10- L.1x<io-6 [Ba] x [SO] =1.2 10-10
PR ge OY isc <8 1.3 7.5X10-3 [Ba] X[F]?=1-7 10-6
2S 6S At lena 1.3X10-? Loxi0-* [Ca] X[CO;] =1.7 X10-8
COsG,04;. 2... 5: 8.0X10-3 6.2x10-5 [Ca] X[C,04] =3.8 X10-9
TOUTS. ds 23. 1.5X<10-! [Ca] X[CrO,.] =2.3 10-2
Sa ae a 1.6X10-? 2.0x10-4 [Ca] X[F]?=3.4x10-1!

GENERAL PRINCIPLES

SOLUBILITIES AND SOLUBILITY PRODUCTS AT ROOM
TEMPERATURE—Continued

Solubility in

Substance. Grams per Liter. Afcles wer Lite: Solubility Products.
a 9 SOR 2.0 1.5X10-2 | [Ca]x[SO.]=2.3 10-4
Cus(CNS)2.....| 5.0*10-4 2.1X10-§ | [Cu] xX[CNS]=1.7x10-"
Cush: so 1.2x10-3 6.0X10-4 | [Cu] x[Cl]=1.410-8
i Se 3.0X10-4 8.0X10-7 | (Cu) x[I]=2.6x10-”
oe ae 8.8X10-% 9.2x10-%3 | [Cu]x[S]=8.5x10-45
a OE AES 3 8.6X10-%8 6.0X10-15 | [Cd] x[S]=3.6x10-%
YO) 5 8 Rs MRCS Nee Carb, Mr Sace ici tu) as [Fe] X[OH]?=1.6<10-"
ae 3.4X10-8 3.9X10-19 | [Fe] x[S]=1.5x10-19
LTO UE S PAROMRNS RAND Pte WR amen Vee ae eR aot [Fe] x[OH]*=1.110-%6
Hg.Bre....<..- 3.9X10-5 6.9X10-8 | [Hg.]x[Br]?=1.3 10-7
PS os See es 3.8X10-4 8.0X10-7 | [Hge] X[Cl]?=2.0 10-18
7 2 Fe eae 2.0X10-7 3.1X10-1° | [Hgo] x[I]?=1.2 10-78
BROCE eee ens. ata eds eee peas [Hg] X[OH]?=4.3 X10- 16?
cS ARR etnies ag CMR en cA aE Rom BR {Hg] x[S]=4.0 10-53
KaPtCle,.: es 11. 2.3xX10-2 | [K]?x[PtCle] =4.9 X10-5
MgCos...'.\... 4.3X10-! 5.1X10-% | [Mg] X[COs] =2.6 10-5
Mg(OH)>.\.: .. 1.2Xx10-2 2.0X10-4 | ([Mg]x[OH]?=3.4x10-
MgNH,PO,,. 8.6X10-3 6.3xX10-5 | [Mg] X[NH4] X[PO,] =2.510-#
DOOR oes Heck een eae ae See [Mn] X[OH]?=4.0 10-4
Muse 2. 38X10 +s 3.8X10-8 | [Mn]x[S]=1.4x10-%
i ee eat ie OKO 1.2X10-12 | [Ni]xX[S]=1.410-%
(ei ag a 9.7 2.7X10-2 | [Pb] X[Br]?=7.9 10-5
Pate aS 11. 3.9X11-2 | [Pb] x[Cl]?=2.4x10-4
PHOO ee 1.1X10-3 4.1X10-§ | [Pb]x[CO,]=1.7x10-"
Phin ees 6.8x10-! 1.5X10-* | | (Pb) xX[f?=1.4x10-t a5
Pb;(PO,)2...... 1.4xX10-4 1.7X10-7 | [Pb]*[PO.J?=1.5 10-3
eS eA eae ae 4.2X10-2 1.5X10-4 | [Pb]x[SO.] =2.3 10-8
PbCrO,....... 4.3X10-5 1.3X10-7 | [Pb] x[CrO.]=1.8 10-4
Mi. OG o28 4.9X10-” 2.0X10-" | [Pb] x[S]=4.2x10-%8
Lt ae 6.6X10-2 3.8X10-4 | [Sr] x[C,0.]=1.410-?
BroUs..fi3.. 1.0x10-? 6.8X10-5 | [Sr}x[CO3] =4.6 10-9
ate Ee tas Ns 2.3X10-1 1.8xX10-% | (Sr]x[F]?=2.5x10-9
BRE os cet 1.1X10-3 6.0X10-4 | [Sr] x[SO.]=3.610-7
FEBS aa sarees 4.8X10-! 1.7X10-% | [Tl]<[Br] =2.9x10-§
oy Cama 3.2 1.2x10-2 | [TI]X[CNS]=1.4x10-4
=: AR Ae oad 3.4 1.4x10-2 | [Tl] x[Cl]=2.010-4
<4 Ra git hor 6.4x10-? 1.9x10-4 | [TI]x{I]=3.6x10-8
MET CS § SERRA ts Ag: Cadet eae ak ay Leah seal Nees [Zn] X[OH]?=1.8X10-"4
aaa Ri TE 3.5X10-12 | [Zn] x[S]=1.210-28

a: 10—%

The above table is prepared from many sources, and the values
are based, in some cases, upon solubility determinations by methods
which are now considered inaccurate.

The table gives a good idea,
however, of the relative order of magnitude.

For copper sulfide,

the table states that 8.8x10-?! gms. dissolve in 1 liter of water.
Obviously, the experimental determination of such a small value is

fraught with difficulty and the probable error is large.

For most

”

Sa sl! oA ae -" a a, AS La es ee ili or ts 5 Sy

SOLUBILITY PRODUCT 23

purposes such a value represents a negligible quantity and the state-
ment is often made that copper sulfide is insoluble in water. It is
instructive, however, to compare the solubility products of the various

sulfides and important methods of separation have been based upon
‘such studies. Only two significant figures have been given in the

table, although it is obvious that more would be justifiable in the
case of the more soluble substances, while even the first figure is doubt-
ful for the very insoluble substances. The values are affected to
different degrees by changes in temperature and the presence of
other substances in solution. A careful, critical study of all the experi-
mental data would be necessary to give the proper number of signifi-
cant figures and it would be necessary to give the exact temperature.

In computing the solubility products, the assumption has been
made that the ionization is complete. Such an assumption is not
permissible with a substance such as ferric hydroxide, and in such
cases only the approximate value of the solubility product is given;
the molar solubility and the grams per liter are not stated. In other
cases the ionization is abnormal as noted.

Two examples will be given to illustrate the method of computing
the molar solubility, S,,, and the solubility product, S,. A saturated
solution of silver iodide contains 3.010-§ gms. (=0.0030 mg.)
per liter. The molecular weight of silver iodide is 234.8. The satu-

3.0X 10-6 | es
S35 = 1-3X10-8. At this

dilution the dissolved silver iodide can be assumed to be completely
ionized:

rated solution, therefore, contains

Agl @ Agt+I-,

and since 1 mole of silver iodide furnishes 1 mole of silver ions and 1
mole of iodine ions, it is evident that the solubility product, S,, is for
Paw (li=[1.3 X10-*| X[1.3 10-8] =1.7 X10-%=§,.

A saturated solution of silver phosphate contains 6.510% gms.
(=6.5 mgs.) per liter. The substance is much more soluble in water
than silver iodide, but its solubility product is smaller. The molecular
weight of silver phosphate is 418.7. The saturated solution, there-

_-3
2X7 =1.6x10-5 moles of silver phosphate which

can be assumed to be completely ionized:
AgsPO4 = 3Agt+PO.,

1 mole of silver phosphate yielding 3 moles of silver and 1 mole of
phosphate ions. The solubility product is

[Ag]? X[PO4] =[3 X1.6 X10-5]* X[1.6 X 10-5] = 1.8 X 10-18 =§,.

fore, contains

24 GENERAL PRINCIPLES

Complex Ions

Silver chloride is slightly soluble in water; 1 liter dissolves about
1.5 mgs. It dissolves very readily in dilute ammonia. The following
reaction takes place:

AgCl+2NH3 @ Ag(NHs3)2Cl.

A study of the properties of this new substance shows that it dis-
sociates in aqueous solution chiefly in this way:

AgNH;Cl 2 Ag(NH3)2*-+Cr-.

The ionic changes involved in the last two equations may be
expressed thus:
Agt+2NHz3 = [Ag(NHs)2]*

and, in accordance with the law of mass action, the greater the con-
centration of the ammonia, the greater the extent to which the reac-
tion takes place in the direction left to right. In a normal solution
of ammonia, the ratio of the concentration of the [Ag(NH3)e] ion to
that of the simple Ag ion is about 107: 1. The [Ag(NHs3)o| ion differs
from the simple Ag ion in much the same manner as the ClO, ClOsz, or
ClO3 ion differs from the simple Cl ion. It is called a complex cation.

When potassium cyanide is added to silver nitrate solution a white
precipitate of silver cyanide is formed:

KCN+AgNOz3 = KNO3+AgCN,
but if an excess of potassium cyanide is used the precipitate dissolves
AgCN+KCN = KAg(CN)z.
In this case the ionic changes may be expressed as follows:
Agt+CN7- @ AgCN,
AgCN+CN™ @ [Ag(CN)e]~

and the silver has become a part of the anion. In this case the value
of the ratio of complex ion to simple ion is even larger than in the case
of the silver ammonia cation.

Similarly, when insoluble ferrous cyanide is treated with an excess
of potassium cyanide, it dissolves, forming potassium ferrocyanide,

Fe(CN)2+4KCN = K4[Fe(CN)ol.

This salt gives none of the ordinary reactions of ferrous ions. The
iron forms an integral part of the complex ferrocyanide ion which has

COMPLEX IONS 25

its own characteristic reactions, and during electrolysis alwavs migrates
toward the anode:

K,[Fe(CN)¢] — 4Kt+Fe(CN),~~.

It is, in fact, quite common to find that simple salts, particularly
in concentrated solutions, are capable of forming such complex com-
pounds. The simple ions can unite with neutral molecules, or with
ions of opposite charge, to form complex ions. If a simple ion adds
to itself a neutral molecule, such as H2O, H2O02, NHs or organic radicals,
then neither the original valence nor the electric charge is changed.
Thus the trivalent cobalt ion is capable of forming a deep red ion with
6 molecules of ammonia and this complex ion has a trivalent charge
like that of the original simple ion.

Cottt+6NH3 @ [Co(NH3)¢]t**.

If, in such a complex ion, one or more of the ammonia groups is
replaced by a negative univalent ion, the valence of the complex ion
is reduced one for each atom of negative ion thus entering into the
complex. :

(NOz)2|*, (NO2)s3

In the presence of potassium cyanide, the trivalent cobalt ion
unites with six CN ions to form a complex which has, in accordance
with the above rule, a triple negative charge. The valence or electric
character of a complex ion is the algebraic sum of the valencies or elec-
tric charges of the constituents.

As already indicated, the stability of these complex ions varies.
When the complex is-very stable, the common reactions of the con-
stituents are not shown. |

Besides these complex ions certain double salts are known. Thus
potassium and aluminium sulfates crystallize together, forming an
alum, KeSO4-Ale(SO4)3-24H20. When this salt is dissolved in
water, the solution shows all the reactions for potassium, aluminium
and sulfate ions, and there is little evidence of the formation of a
complex ion. To determine, in a given case, whether a substance
is a double salt or a complex salt, it is customary merely to see whether
the characteristic reactions of the simple ions are shown. A salt
exists which has the symbol KCr(C204)2:5H20. An aqueous solu-
tion of this salt readily shows the reactions for the potassium ions,
but reacts sluggishly when tested for chromium cations or for oxalate
anions. Evidently the chromium and the oxalate have united to form

26 GENERAL PRINCIPLES

a complex anion with a negative valence of one, but this complex is
not as stable as some of the others that have been mentioned. It is
probable that there is no sharp distinction between the double salts
and the complex salts- and probably the double salts are most logic-
ally to be classed as complex salts of which the complex ion is not
very stable. As a general rule, those complex salts which are composed
of neutral salts of strong acids yield complex anions which are largely
dissociated into simple ions in dilute solution. On the other hand,
complex ions composed of positive ions and anions of weak acids are
usually very stable.

Reactions of the Ions

As already indicated, most of the reactions used in qualitative
analysis involve reactions between ions. We have seen.that, in prin-
ciple, all reactions are reversible and have learned to understand
some of the laws which govern these reversible reactions. In analyti-
cal chemistry, it is necessary for the most part to employ reactions
which take place almost completely in the desired direction. Unless
a reaction can be made to go nearly to completion in a given direction,
it is of little value either as a sensitive test or for furnishing a method
of separation. The useful reactions of qualitative analysis, namely
those which-apparently go to completion, may be brought into four
classes: :

(1) Reactions in which a gas is formed.

(2) Reactions in which a precipitate is formed.

(3) Reactions in which a non-ionized substance is formed.

(4) Reactions of oxidation and reduction.

When a gas is formed as a result of a chemical reaction and the
gas escapes, the reaction will go to completion. All gases can be
boiled out of solution and thus all reactions of this type can be made
to go to completion. The reaction can be stopped by preventing
the escape of the gas; this shows that the reaction is inherently a
reversible one.

Whenever a substance which has a very small solubility product
is formed by means of a chemical reaction, the greater part of the sub-
stance will leave the solution in the form of a precipitate and the
reaction will go practically to completion. The table on page 21
shows that the saturated solution of silver chloride contains only
about one hundred-thousandth of a mole (=0.01 millimole) of solid
salt per liter. The table also shows that when the product obtained
by multiplying the concentration of the silver ions by the product
of the concentration of the chlorine ions in any aqueous solution is

OXIDATION AND REDUCTION 27

equal to 1.3X10-!° the solution is saturated with silver chloride.
By adding an excess of chlorine ions to a solution containing silver ions,
it is possible, therefore, to precipitate nearly all of the silver. It is
evident that it will take less silver ions to give the solubility product
when an excess of chlorine ions is used than is necessary when pure
silver chloride is dissolved in water.

The precipitated silver chloride will dissolve completely in potas-
sium cyanide, because the silver ion forms with the cyanide ion a
complex which is ionized to such a slight extent that the solubility
product of silver chloride is no longer reached, even although all the
chlorine is present in the ionic condition.

The formation of a non-ionized substance also causes a reaction to
go to completion. The table on page 10 shows the ionization values
of a few common substances. This table may be used exactly like that
of the solubility products to enable one to predict whether a reaction
is likely to go in a given direction. The equilibrium between water
and its ions H and OH has been discussed on page 6. The same
reasoning may be applied to the equilibrium between any other slightly
ionized substance and its ions; whenever the ions are added sepa-
rately to a solution, some of the non-ionized substance is at once
formed. ‘Thus when any acid is added to the solution of a sulfide
a reaction takes place, partly because the hydrogen sulfide is a very
weak electrolyte and partly because the substance is a gas. Sim-
ilarly calcium phosphate dissolves in hydrochloric acid because
more PO, ions are formed from dissolved calcium phosphate than
are formed from HzPO,4~ ions in the presence of an excess of H* ions
from the hydrochloric acid; the reaction takes place because of the
formation of a non-ionized substance.

Finally, many reactions of oxidation and reduction take place nearly
to absolute completion, although all these reactions can be shown ~
to be inherently of a reversible type. To understand such equilibria,
however, it is necessary to discuss oxidation and reduction at greater
length.

Oxidation and Reduction

The term oxidation, in its narrowest sense, signifies the taking
up of oxygen by an element or compound. Thus ferrous oxide, on
being heated in the air, is converted into ferric oxide and the reaction
is called an oxidation. Since, however, ferric chloride bears the same
relation to ferrous chloride that ferric oxide bears to ferrous oxide,
it is customary to call the change of ferrous chloride into ferric chloride
an oxidation, although it is not necessary to think that oxygen takes

28 GENERAL PRINCIPLES

part in the reaction at all. This is an interesting example of a word
in common use which has come to mean a great deal more than it
originally meant. Indeed, chemists have departed so far from the
original meaning of oxidation that sometimes the word seems inappro-
priate, and the use of another word, such as adduction, has been sug-
gested. Reduction is the exact opposite to oxidation, and whenever
one substance is oxidized some other substance is reduced. Hydrogen
was formerly considered to be the typical reducing agent, so that the
definition for oxidation used to read something like this: Oxidation
is the addition of oxygen (or its equivalent) to an element or com-
pound or the taking away of hydrogen (or its equivalent).
The reaction between ferrous chloride and chlorine:

2FeClo+Cle = 2FeCls,
expressed in terms of the ionic theory becomes,
| 2Fet+++Cle=2Fet + ++42Cl-

In other words, the diferrion has been converted to triferrion and the
neutral chlorine molecule has become changed to negatively charged
chloride ions. In all other reactions in which a ferrous salt is oxi-
dized, the valence of the iron is increased one, and the modern concep-
tion of oxidation and reduction is summed up very simply as follows:

Oxidation is the increase in the valence of an element or radical in the
positive direction; reduction is the increase in the valence of an element
or radical in the negative direction. Oxidation involves the assumption
of positive charges or the loss of negative charges and reduction in-
volves the loss of positive charges or the assumption of negative charges.

According to the electronic conception of the constitution of matter,
the atom of an element consists of positively charged corpuscles and
negatively charged corpuscles or electrons. The mass associated
with the positive electricity is much larger than the mass associated
with the equal charge of negative electricity. The number of positive
and negative electrons which make up the atom is probably a very
small multiple of the atomic weight of the element. ‘The mass asso-
ciated with a unit negative charge is so small that it may easily be lost,
but only under conditions such that it is accepted by some other atom.
The originally neutral atom which loses the electron thus becomes
positively charged and the atom which accepts the negatively charged
electron becomes negatively charged, and a tube of force holds the
two elements together in a so-called chemical compound. In the light
of the electron theory, therefore, an element is oxidized when it loses
an electron and an element is reduced when it receives an electron.

OXIDATION AND REDUCTION 29

This is the simplest, and at the same time most comprehensive theory
of oxidation that has ever been suggested. |

‘Oxidation, according to this conception, is essentially an electric
phenomenon. This theory suggests the thought that it ought to
be possible to accomplish oxidation and reduction simply by means of
electric energy. As a matter of fact probably every oxidation and —
reduction can be brought about in the electrolytic cell if the proper
conditions be maintained. Using the conventional symbol @ to
designate a unit charge of positive electricity and ©) to designate a
unit charge of negative electricity, but bearing in mind that the nega-
tive electricity is alone transferred and that the only way an element
can gain in positive charge is by losing one or more negative electrons,
we may express oxidations in the electrolytic cell as follows:

Fett+ @—Fettt or Fet*—©— Fet tt.

Such oxidations take place at the electrode called the anode. Con-
versely, at the cathode, ferric salts can be reduced to the ferrous
condition:

Fet*+++@©-— Fett.

_ Not only may all oxidations and reductions be accomplished with
the aid of the electric current, but, vice versa, an electric current may
be produced by a proper arrangement of the components of any reac-
tion of oxidation and reduction. Thus some ferric chloride and sodium
chloride solution in a small beaker may be connected with a second
beaker containing sodium chloride by means of a U-tube filled with
dilute salt solution. If a platinum electrode is placed in each beaker
and the terminals are connected with a sensitive voltmeter, no cur-
rent will pass through the wire. On pouring some hydrogen sul-
fide water into the beaker containing sodium chloride, a decided deflec-
tion of the voltmeter needle is at once observed, showing the passage
of an electric current. The negative current enters the voltmeter
from the solution containing the hydrogen sulfide and passes on to
the ferric chloride solution and back, through the salt-bridge, to the
hydrogen sulfide solution. The chemical reaction that takes place is,

2Fet t+49= — 2Fett+8.

The electric current is produced as a result of the oxidation of the
sulfide ions and reduction of the ferric ions.

Oxidation and reduction reactions are inherently reversible reac-
tions, like all other chemical reactions, and are effected by the concen-

30 GENERAL PRINCIPLES

trations of the reacting substances. Thus, in the above experiment
the intensity of the electric current can be greatly increased by using
a soluble sulfide instead of hydrogen sulfide, the former being more
largely dissociated and yielding a larger concentration of sulfide ions.
Or, by adding a fluoride to the solution of ferric chloride, a fairly stable
complex ion, Fel’, is formed and the current slackens, owing to the
decreased concentration of the ferric ions.

The most important oxidizing agents used in analytical chem-
istry are the halogens, nitric acid, potassium permanganate, potassium
dichromate and hydrogen peroxide.

The most important reducing agents are nascent hydrogen and
metals, sulfurous acid, hydrogen sulfide, stannous chloride and hydri-
odic acid.

The oxidizing action of halogens depends upon the conversion
of the neutral halogen into halogen anions.

The oxidizing action of halogen upon ferrous ions results in the ©
formation of ferric ions and of halide ions:

2Fet*++Cle > 2Fettt+2Cr,
or
2FeCle+Cle = 2FeCls.

The action of halogen upon hydrogen sulfide is interesting. First
of all, the sulfid-ion is oxidized to free sulfur,

H2S+Bre — 2HBr+5,

but, if the bromine is fairly concentrated, the reaction may go farther
and the sulfur be converted into sulfuric acid, the whole reaction being

H28+4Bre+4H20 = H2804+8H Br.

It will be noticed that it is very easy to balance equations of
oxidation and reduction by noticing the change in valence. In this
last equation sulfur goes from a negative valence of two to a positive
valence of six, making an algebraic change of eight, which corresponds
to the loss of eight electrons by the sulfur atom.

The oxidizing action of nitric acid depends upon the reduction
of the nitrogen. The extent of the reduction depends upon the con-
centration of the nitric acid and the nature of the substance oxidized.
The more concentrated the nitric acid, the less it is reduced; the more
concentrated the reducing agent and the stronger its reducing power
the greater the reduction of the nitric acid. Nitric oxide, NO, is
commonly formed, but often other products such as nitrogen peroxide,
nitrous oxide, nitrogen and even ammonia are produced.

OXIDATION AND REDUCTION 31
\

In nitric acid, the nitrogen atom has five positive electric charges
residing upon it. When it is reduced to nitric oxide, NO, it has only
two positive charges, the nitrogen having accepted three electrons.
The reaction between a ferrous salt and nitric acid is,

3Fet*++NO37+4H*3F et +*+NO+2H20,
or
6FeSO4 + 2HNO3 +3H2S04 =3Fee (SO) 3+ 2NO +4H20.

The reaction is doubled in the last instance simply to get an even
number of molecules of Fee2(SO4)3. The addition of sulfuric acid is
unnecessary, as the nitric acid can also act as an acid, in which case
a mixture of ferric sulfate and nitrate is formed:

3FeSO4+4HNOsz = Fe2(SO4)3-+Fe(NO3z)3 +-NO+2H20.

The action of nitric acid on a sulfide is interesting. If the nitric
acid solution is cold and dilute (0.8 N) there is hardly any oxidation
of the sulfur:

MnS+2HNO3 = Mn (NOs)2 +H>S.

If the nitric acid is more concentrated (e.g., 2 N) and the solu-
tion is heated, the sulfide is changed to nitrate and free sulfur is
formed. Thus for the reaction between copper sulfide and hot nitric
acid, each atom of copper requires two nitrate ions, each atom of sul-
fur loses two electrons and, in accomplishing the oxidation, each atom
of nitrogen gains three electrons. The reaction may be expressed
thus:

3CuS +8HNO3 =3Cu(NO3)2+4H20+35+2NO,
or
3CuS+8H*t+2NO3— — 3Cut*+4H20+38S+2NO.

If the nitric acid is very concentrated, the greater part of it is
reduced only to NOzg and the sulfur is oxidized to sulfuric acid. The
reaction may then be written:

CuS+8HNO3=CuSO04+8N02+4H20.

The oxidizing action of permanganate depends upon the readi-
ness with which the manganese is converted into a manganese com-
pound of lower valence. In the permanganate anion, MnO, the
manganese has a positive valence of seven. Ordinarily, in acid solu-
tion, the permanganate is reduced to bivalent manganese cation,

32 GENERAL PRINCIPLES

corresponding to a loss of five positive charges, or acceptance of five
electrons:
MnO4~+5Fet*++8H*=Mnt*t-+5Fet*+t+4H.20;

2MnO4 +5H2S+6H* =2Mn*t *+5S+8H20;

++
2Mn0O4~+5S8n* sae 16Ht a 2Mn* F4 58nt *+8H2O;
2Mn0O4 +10HI+ 6Ht = 2Mntt-+5lo +8Ho20.

The oxidizing action of a chromate or dichromate ordinarily
depends upon the formation of trivalent chromic ions. In the chro-
mate and dichromate ions the chromium atom has a positive valence
of six, so that for each atom of chromium the reduction corresponds to
a loss of three positive charges, or gain of three electrons. Potassium
chromate in acid solution is in equilibrium with the dichromate:

9CrO4-+2Ht — 2HCrO4~ — H2O+Cr207-.

In balancing equations, therefore, it makes little difference whether
we start with the chromate or dichromate, except with respect to
the quantity of acid required:

Cr207~+6Fet*+14Ht — 2Crt*+*++6Fet **++7H20;
Cre07~+38"+14H* > 2Cr++*++38+7H20; |

+4
Cro077+38n* ++14H*t > 2Crt*+*+3Snt *+7H20;
CreO7_ +61 + 14H*t — 2Cr+*+*+31.+-7H20.

Hydrogen peroxide acts both as an oxidizing agent and as a
reducing agent. It oxidizes ferrous chloride to ferric chloride and
it is capable of reducing permanganate to manganous salt. This
anomalous behavior has been the cause of considerable discussion
in the literature. It is unnecessary to go into the details of such a
discussion, but a simple explanation of this behavior will be suggested.
In all the other compounds of hydrogen and oxygen that we shall
study, hydrogen has a positive valence of one and oxygen a negative
valence of two. With hydrogen peroxide, also, it is best to assume
that the hydrogen has its normal valence corresponding to one posi-
tive charge. ‘Two structural formulas are at once suggested for hydro-
gen peroxide,

H—-O
OO or Pe
H—O

Hy,
H

OXIDATION AND REDUCTION 33

Each of these formulas gives to one atom of oxygen its normal negative
charge of two units of electricity and each gives to one atom of oxygen
an equal number of positive and negative charges. It is unnecessary,
therefore, to attempt to decide which of these formulas is the more
appropriate.

The characteristic behavior of hydrogen peroxide may be traced

to the presence of the atom of oxygen which has an equal number of

positive and negative charges. In alkaline solution, hydrogen
peroxide decomposes spontaneously and oxygen is evolved. This
spontaneous decomposition, with the formation of neutral oxygen, is
easy to understand on the basis of the assumption that the original
molecule already contains the atom of oxygen in a very unstable con-

dition of neutrality.

The oxidizing power of hydrogen peroxide is due to this atom of
oxygen. In acid solutions ferrous iron is converted by it to the ferric
condition. The peculiar atom of oxygen loses its positive charge and
receives in its place a negative charge; the total change corresponding
to the acceptance of two electrons:

2Fet ++-H20.+2H+t > 2Fe++++2H.20.

The reducing power of hydrogen peroxide also depends upon the
presence of this atom of oxygen. When in contact with a powerful
oxidizing agent, such as permanganate or another peroxide; a reac-
tion takes place and oxygen gas is evolved. It has always been
assumed, since the classic experiments of Schénbein, that half of the
evolved oxygen comes from the oxidizer and half from the hydrogen
peroxide. It is simplest to assume, therefore, that the oxygen is
momentarily changed to an atom with two positive charges which at
once unites with negatively charged oxygen in the oxidizer; or, this
oxygen in the hydrogen peroxide may unite with similarly charged
oxygen in another peroxide.

2MnO4 +5H20e2 +6Ht —? 2Mn* *4+8H20+50p;
Mn0O2+H202+2H* — Mnt *42H20+0Og;
Coo0s +HeOe +4H* ae 2Cot +43H.O +Ood.

The characteristic action of the more important oxidizing agents
has now been considered briefly and it remains to describe the char-
acteristic behavior of the important reducing agents. Since every
oxidation involves a simultaneous reduction, all of the above reac-
tions can be used to illustrate reduction as well as oxidation.

The reducing action of nascent hydrogen and of metals depends
upon the conversion of the neutral hydrogen or metal into positively

34 GENERAL PRINCIPLES

charged cations. Such a reduction may take place in acid, alkaline, —
or neutral solution.
(a) In acid solution, by the-employment of zinc, ete.: -

Zn+H2S804= ZnSO4+ He
or
Zn-+2H*t — Zn* t+Hbp.

This reaction in itself represents both an oxidation and a reduction,
inasmuch as the,metallic zinc, which is electrically neutral, becomes
changed into zinc with two positive charges and, on the other hand,
the hydrogen in sulfuric acid has lost two charges and become electric-
ally neutral hydrogen.

Now this nascent hydrogen, as fast as it is formed, may act as a
reducing agent and serve for effecting the reduction of some other
element, e.g., ferric chloride; in which case the final changes are the
oxidation of the zinc and the reduction of the iron:

Fet t*+4-7n=Zn' T+-2F ett.

In such cases it is doubtful whether the reduction process really goes

through the stage of forming nascent. hydrogen. :
By the action of zine and very dilute sulfuric acid, it is easy to

transform silver chloride into metallic silver:

2Agt+Zn — Znt++2Ag,
or an arsenite into arsine:

AsO3- +3Zn+9H* — 3Zn**+3H20-+-AsHs.

The result of this last equation is the oxidation of zine from the metallic
condition to the bivalent state and the reduction of the arsenic, which
is given three negative charges in place of the three positive valences
that each arsenic atom has in AsOg3.. Thus each As atom loses six
charges of electricity and each Zn atom gains two, so that it takes
three Zn atoms to reduce one AsO3 anion.

(b) In alkaline solution, by means of zinc, aluminium, sodium
amalgam, or by Devarda’s Alloy (Cu=50, Zn=5, Al=45). This
reaction also is sometimes attributed to nascent hydrogen: |

Zn+2Na0H = NaeZnOe+He or Zn+20H — ZnO2 +He..
2Al+2Na0H+2H20 =2N ahOe st Wile or 2Al+20H~ +2Hs0O
> 2A1097++-3Hpe.
In the case of Devarda’s Alloy, the reaction is completed mudh’tore
quickly than by the use of either zine or aluminium alone. Nitrates

ee ee ee
:
“

OXIDATION AND REDUCTION 35

and chlorates may be reduced in a few minutes by means of Devarda’s

_ Alloy and a few drops of sodium hydroxide; the reaction also takes:

place in neutral solution, but it takes considerably longer:
. NO3 +4Zn+70H~ — 4Zn0O2~+2H20+NH3
ClO3"+3Zn+60H- — 3Zn02*+3H20+CI.

In the nitrate-ion nitrogen has a positive valence of five; in
ammonia it has a negative valence of three. By the reduction with
zinc, therefore, the nitrogen loses eight positive charges, or accepts
eight electrons. At the same time the zine accepts two electrons,
forming, in a neutral or alkaline solution, the zincate-anion. Thus

one atom of nitrogen in the nitrate-ion requires four atoms of zine to

convert it into ammonia.

Similarly, the chlorine atom in the chlorate-ion has a positive valence
of five and is reduced to a negative valence of one by the reaction with
zine in neutral or alkaline solution. Thus one chlorate-ion reacts with
three atoms of zinc. Inspection of the above equilibrium expression
shows that one could predict that the reaction would take place best
in alkaline solution in accordance with the mass action law. ,

Reduction by means of sulfurous acid takes place in moderately
acid solution and depends upon the fact that sulfur is more stable
when it has six positive charges, as in sulfuric acid, than when it has
only four as in sulfurous acid. Ferric salts are readily reduced by
this reagent, and since the iron loses only one valence while the sul-
fur gains two, it is evident that one molecule of sulfur dioxide (the
anhydride of sulfurous acid) will reduce two atoms of iron in a ferric
salt:
Fe2(SO4)3 +2H20+S02 = 2H2804+2FeSOz,
or

o¥et * *+-S0."+-Hs0 — 2Fe* *-++-S0.7°+-2H".

In a similar manner, the arsenate-ion and many other substances
are reduced very readily by means of SO2 or SO3 ':

AsOaz +S037>—AsO3_ +5804-

An excess of aqueous sulfurous acid is added to the solution which is
to be reduced; it is then heated to boiling; and the boiling is contin-
ued while a stream of carbonic acid gas is passed through the solu-
tion until the excess of sulfurous acid is driven off.

Reduction by means of hydrogen sulfide, in which the sulfur
atom possesses two negative charges, depends upon its oxidation to
free sulfur, which is electrically neutral. Thus two atoms of ferric iron

36 . GENERAL PRINCIPLES

are reduced to ferrous iron by one molecule of hydrogen sulfide and

one molecule of potassium dichromate reacts with three molecules of

hydrogen sulfide: | |
2Fet*+++HS — 2Fett+2Ht+S

Cr207-+3H2S+8H* — 2Cr++++3S+7H20.

One objection to the use of hydrogen sulfide as a reducing agent is
the difficulty involved in the subsequent removal of the precipitated
sulfur by filtration. Moreover, hydrogen sulfide is used in qualitative
analysis chiefly as a precipitant. If a solution contains an oxidizing
agent (such as nitric acid, chloric acid, chromic acid, etc.), the sulfide-
ion will be oxidized and there will be separation of sulfur. Any
_ sulfide obtained will be largely contaminated with sulfur, which renders
the subsequent examination more difficult. If the solution contains no
metal which is precipitated by hydrogen sulfide, but contains oxidizing
agents, it will still cause separation of sulfur. One is often in doubt
whether there is not some sulfide mixed with the sulfur, and is therefore
obliged to examine the precipitate further, which is often unnecessary if
the oxidizing agent is previously destroyed. Hydrogen sulfide reduces

Halogens: H28+Cle 3 =2HCI+S;

Nitric Acid: 2HNO3+3H28 =4H,0+2NO0+38;
Chloric Acid: HC10O3+3H28 =3H20+HCI1+35;
Ferric Salts: 2FeCls-+ H2S =2HC1+2FeClo-+S;

Chromic Acid: . 2Cr03+3H2S+6Ht =6H20+2Cr*++t+38;

Permanganic Acid: 2HMn04+5H28+4H* =8H20+2Mn**+58;
and many other substances. .

Reduction with stannous chloride takes place usually in acid
solutions. The reduction depends upon the fact that stannous ions
are readily changed to stannic ions: |

SnCle +Clo = SsnCh,
or +40
Snt*++Cle=Snt*t+2Cr,
Ferric salts, chromates, permanganates, mercuric salts, and many
others are reduced in this way:

hed '
2Fet*+++S8ntt > 2Fett+Sntt;
++
2CrO4~+3S8n*t ++16H* > 2Crt*+*+3Snt *+8H20;
+4
2HgCle +8nt +t — Sntt+t+HgeCle+2Cr; |

seer iy
Hg2Cle+Sn** > Sn**+2Hg+2Cr-.

ee ee ee

- ae 22 Pat

a ue Je <<? e
F 3 2

ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 37

Reduction with hydriodic acid depends upon the change of the

iodine anion into free iodine. Most substances that are capable of

being oxidized or reduced readily can be made to react with either
hydriodic acid or with free iodine. It is easy to detect the presence of
free iodine, and for this reason the iodometric reactions are extremely
important in the study of analytical chemistry. To prevent the
oxidizing effect of free iodine, an excess of potassium iodide is usually
required, and means are often taken to remove the iodine as fast as it
is formed; this is in accordance with the mass-action principle,

2Mn04-+101-+16H* > 2Mn*++8H20+5lo;
Cr207~+61 +14Ht > 2Cr*+*+*++7H20+3!I2;
Ket? +1 => Pet t+I.

Electromotive Series and Oxidation Potentials

If a substance like sugar lies as a solid on the bottom*of a beaker |
filled with water, the molecules of sugar tend to distribute themselves
throughout the solution; in other words the sugar dissolves. The
tendency of the solid molecules to pass into solution may be regarded
as the result of pressure and, in fact, it is customary to say that the
solid substance possesses a solution pressure.

If sufficient solid is present, eventually, with the aid of diffusion,
the liquid will reach a state of saturation. The liquid then contains
an equal quantity of sugar in all its parts and, at the prevailing temper-
ature, will not dissolve any more sugar. There must, therefore, be
some force which acts in opposition to the solution pressure and pre-
vents a saturated ‘solution from dissolving any more of the solid
substance. This force is the osmotic pressure which the dissolved mole-
cules exert in the solution. In a saturated solution, the osmotic pres-
sure, which is itself determined solely by the number of molecules of
dissolved substance and the temperature, exactly balances the solu-
tion pressure of the solid substance. The process of dissolving a solid
substance involves no electrical effects. This is also true when the
dissolved substance is an ordinary electrolyte, because an equal number
of positive and negative ions is formed and there is no electric dis-
turbance.

The metals themselves, though usually to a much less degree, also
show a tendency to dissolve when placed in contact with water. In
this case, however, an oxidation takes place, for, to the extent that it
dissolves, the metal is converted into electrically chargedions. The

38 GENERAL PRINCIPLES

tendency of the metal to dissolve is called its electrolytic solution pres-
sure. Just as in the case of the sugar, the osmotic pressure of the ©
dissolved ion acts against the solution pressure. The electrolytic
solution pressure has a definite value which is characteristic of each
metal.

' If a metal such as zinc, which oxidizes fairly readily, is placed
in a saturated solution of zinc sulfate, none of the metal dissolves. If
it is placed in contact with a dilute solution of zine sulfate, the solu-
tion pressure of the zine is greater than the deposition pressure and
some positively charged zine ions pass into solution. Thereby, the
metal itself acquires a negative charge and the solution a positive
charge. Asa result of the charge residing upon the zinc ions that have
gone into solution, an electrostatic force is produced which seeks to
force the ions back upon the metal. This electromotive force is added
to the osmotic pressure of all the zinc ions in solution and it increases
rapidly with the number of ions that dissolve from the metal. When-
the sum of the osmotic pressure plus the electromotive force is equal
to the electrolytic solution pressure of the zinc, the zine stops dissolving.

When a less-readily oxidizable metal, such as copper, is placed in a
copper sulfate solution the relations are reversed. In this case, except
in extremely dilute solutions, the osmotic pressure is. greater than
the solution pressure and the metal does not dissolve; on the contrary,
a few of the copper ions are discharged on the metal, giving to it a posi- —
tive charge while the solution becomes negatively charged. Equilib-
rium is established as soon as a few of the ions have been thus deposited.
The potential difference between the metal and its solution, or as it is
often called, the potential of the metal, is said to be positive when the
charge of the solution is positive; this is the case with the readily oxidiz-
able metals such as magnesium, aluminium, zinc, iron, etc. On the
other hand, the potential of the metal is negative when it is difficultly
oxidizable; this is the case with copper, silver, platinum and gold. —

Simple contact of a metal with a solution of its ions usually
results in a potential difference between the metal and the solu-
tion. Such a potential is determined by the relation that exists
between the electrolytic solution pressure of the metal and the
osmotic pressure of the solution. Equilibrium is soon reached in
most cases and the simple contact of a metal with a solution of its
ions is not a permanent source of electricity. If, however, two metals
of different potential are placed in contact with their respective solu-
tions, then electric charges of different potentials result, and if the two
metals are connected outside the liquids by a wire, an electric current
flows from the higher potential to the lower. Since the original differ-

‘

ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 39

ences in potential between the solutions and the metals are constantly
being reéstablished, a permanent current results. This is the prin-
ciple of the Daniell cell, in which a normal solution of copper sulfate.
is separated by a porous partition from a normal solution of zine |
sulfate. A zinc rod is placed in the zinc sulfate solution and a copper

~ plate in the copper sulfate solution; the current flows through the wire

from the copper to the zinc and through the solution from the zine
to the copper.

Nernst, who was the first to suggest the above explanation of the
origin of the electromotive force on the basis of the relations of osmotic
pressure, has worked out a formula for computing the potential differ-
ence which exists at the place of contact of a metal with a solution of
its ions. If H denotes this potential in volts, R the gas constant
expressed in voltsXcoulombs, F the electrochemical equivalent or
quantity of electricity borne by one equivalent weight in grams of the
ions of any metal, n the valence of the ions, P the electrolytic solu-
tion pressure, p the osmotic pressure, and 7 the absolute temperature
of the solution, the Nernst formula reads:

eee
Bey Be

Substituting the numerical values for R(8.32) and F(96,500), dividing
by 0.434 in order to use common logarithms, and assuming the ordinary.
room temperature to be 18° C. (=291° absolute), the formula becomes:

ata log 2 volts.

Ege =

Inasmuch as the osmotic pressure, p, depends solely upon the concen-
tration of the solution and the temperature, ,the electromotive force
resulting by the contact of a metal with its ions is shown by the
formula to increase as the electrolytic solution pressure of the metal
increases and to decrease asthe concentration of the ions increases.
Since multiplying a number by 10 simply raises its common logarithm
one whole unit, the formula also shows that increasing the concentra-
tion of the solution tenfold lowers the electromotive force in question
about 0.06 volt when the ions are univalent, 0.03 volt when the ions
are bivalent, and 0.02 volt when the ions are trivalent. Lowering the
concentration until it is one-tenth of its original value raises the elec-
tromotive force nearly 0.06 volt in the case of univalent ions, 0.03
volt with bivalent ions, and 0.02 volt with trivalent ions. The electro-
lytic solution pressure is a measure of the readiness with which a metal
can be converted into its ions; or, since the formation of the ions

40 GENERAL PRINCIPLES

involves an oxidation, it determines the readiness with which the ele-
ment undergoes oxidation. The electromotive force that results can
be appropriately called the oxidation potential. It is also a measure
of the force required to deposit a metal from solution by means of the
electric current.

It is now easy to understand what happens when a metal is placed
in a solution containing the ions of some other metal. It is a well-
known fact that the immersion of a strip of iron in a solution of copper
sulfate causes the deposition of metallic copper, while an equivalent
quantity of iron dissolves as ferrous sulfate. The copper is reduced
to the metallic condition by means of metallic iron and the latter is
oxidized by means of cupric ions. This is because the electrolytic
solution pressure of iron is so much greater than that of copper that a
condition of equilibrium is not reached until practically all of the copper
has been precipitated. The oxidation potential of metallic iron against
a molar solution of a ferrous salt is 0.43 volt and of metallic copper
against a molar solution of cupric ions is about —0.34; the minus sign
means merely that there is more tendency for copper-ions to be de-
posited than for metallic copper to pass into solution. The greater the
positive value of the oxidation potential, the greater the electrolytic
solution pressure. As the copper is deposited from the solution, its
oxidation potential becomes gradually larger, and as the iron passes
into solution its oxidation potential becomes smaller and smaller. The
Nernst formula shows that the value for metallic copper against a
tenth-molar solution of its ions will be raised to —0.31 volt and
the value for iron against a solution of tenth-molar ferrous salt will
be 0.40 volt. It is evident that equilibrium between the iron and the
cupric solution will be reached only when the oxidation potential of
the iron is equal to the oxidation potential of the copper; before this
happens either the solution will become saturated with ferrous sulfate
or all but an infinitesimal quantity of the copper will be precipitated.

It is possible to arrange all the metals in a series according to their
electrolytic solution pressures. Such a series is called the electro-
motive series of the metals. It enables one to understand all the
reactions in which a free metal is involved either as an initial or
a final product. The entire chemical activity of the metals corre-
sponds fairly closely with such an arrangement. The members at. the
top of the series are the most readily oxidizable; those following cop-
per do not oxidize or rust when exposed to the air.

The electromotive series shows the relative value of the metals as
reducing agents. The metals at the top of the series are the best reduc-
ing agents. Thus the alkalies are such good reducing agents that

ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 41

they will even decompose water at ordinary temperatures, reducing
the positively charged hydrogen to the neutral condition.

It is important to remember, however, that it is not alone the electro-
lytic solution tension which determines the oxidation potential. The
concentration of the solution also comes into consideration. If the
oxidation potentials are all measured against equivalent concentra-
tions, then the order of the metals arranged in the electromotive series
will correspond exactly to the order of the metals when

free hydroxyl ions in solution and these are in equi-
librium with the zinc ions.

arranged according to the values of their electrolytic ~t®cTRomotive
: SERIES OF THE
solution pressures. Mineasay
‘ We are now able to understand why the alkalies Cashier
decompose water readily and why the quantity of Rubidium
zinc ions formed under similar conditions is very Potassium
small. The oxidation potential of zinc against a molar spgscadl
~ solution of zinc ions is about 0.76 volt. Water, however, Aluminium
is ionized very slightly; the table on page 10 gives Manganese
its ionization as 2X10-’ per cent. The oxidation re
potential of hydrogen against such a very dilute solu- Thallium
tion of hydrogen is not 0.0, as given in the table, but Iron
it is nearer the zinc value. If the ionization of water hae
were 10-*6, the value would be approximately that of Tin
zinc. According to the oxidation potentials, therefore, Lead
we should expect zinc to decompose water with libera- ae
tion of gaseous hydrogen. As a matter of fact zinc is Bismuth
oxidized somewhat by contact with water and the Arsenic
oxidizing agent is the hydrogen of water; but the ee
reaction does not take place to any extent. The Silver
primary products of the reactions are zinc ions and Palladium
free hydrogen, but the escape of the hydrogen leaves Aes

The table on page 22 gives the solu-

bility product of zinc hydroxide as 1.8X10-'*. The reason the zine
does not decompose water, therefore, is because it is protected by
the film of insoluble oxide or hydroxide which quickly forms upon
it. On the other hand, when the hydrogen is present in the form of
an acid, with the anion of which zinc forms a fairly soluble salt, the
oxidation of the zinc ordinarily continues at the expense of hydrogen
ions until all of the zinc is dissolved.

The electromotive series shows the oxidation tendencies of the
elements, but in attempting to predict what will happen in any given
case it is necessary to bear in mind that the concentration of the solu-
tion must be considered and the solubility relations.

42 _ GENERAL PRINCIPLES

We have seen that the Daniell cell* is obtained by taking advantage
of a difference in oxidation potentials, and it was stated on page 29
that a similar cell could be formed by taking advantage of any reac-
tion of oxidation and reduction. All reactions of oxidation and reduc-
tion take place because of differences in oxidation potentials. Just
as the electromotive series of the metals helps one to predict whether
a metal will act as a reducing agent or not, so a complete table of oxida-
tion potentials will help one to determine whether any given reaction
of oxidation and reduction may be expected to take place in the desired
direction. Such a table of oxidation potentials is given on page 43.
The table shows the values referred to molar solutions, the value of
normal hydrogen ions in the normal hydrogen electrode being taken
as 0. The positive value of the electromotive force shows that the
oxidation takes place very readily, the unchanged substance assuming
a negative charge; or, in other words, if a cell is constructed with
the normal hydrogen electrode, the direction of the current (positive
to negative) in the solution is toward the hydrogen electrode when
the element has a positive potential. The first column in the table
gives the original state of the element or ion, the second column shows
the change in charge that this element or ion undergoes, the third column
the oxidized condition and the fourth column the oxidation potential.

The electromotive force of any reaction of oxidation and reduction
is determined by the difference in the oxidation potentials. Thus
in the Daniell see using molar solutions of copper sulfate and zine
sulfate, the electromotive force of the entire cell is the difference between
the oxidation potential of copper (—0.34 volt) and zine (+0.76) =1.10
volts.

*Jn the well-known Daniell cell the zinc, which is in contact with zine sulfate
solution, is the negative electrode, and the copper, which is in contact with
copper sulfate solution, is the positive electrode; zine passes into solution at one
electrode and copper is deposited at the other. Outside the cell, the positive-to-
negative direction is from copper to zinc, but inside the cell it is from zine to
copper. In measuring the single electrode-potentials of such a cell, all German
and many English texts assign a negative value to the potential of the element
which dissolves and a positive value to the element which is deposited. This is,
however, a purely arbitrary practice, and it seems more natural to assign the
positive and negative values in the opposite manner. It is rational, in the Daniell
cell, to assign a positive value to the electromotive force which measures the
tendency of zinc to form zinc ions, and a negative value to the tendency of cop-
per to form copper ions. In tracing the course of the current in the Daniell cell,
it is simplest to start with the zine which forms zinc ions and pass through the
solution to the copper and back through the wire to the zinc. In the cell, the
zinc is positive with respect to the copper but in the wire the copper is positive
to the zinc.

ELECTROMOTIVE SERIES AND OXIDATION POTENTIALS 43

OXIDATION POTENTIALS

Ag Ag
Original | #8| Higher | E.mf. Original “| 32 | © Higher | Bmt.
tate ES State. in volts. State. ES State. in volts,
=| =|
| SS ae Bet Bava ce +3.03]| Cu®......... a PCr. 52 3: --0.51
: K° oc, oe AHS 1 Kt rahe eit a a +2 “ 93 Mn0O.+ 3 MnO;- +Ht —0 - 52
Pea... ¥( Nee Sc. +2.72 20H— -
Ga aa SV Raw. s: et he, Te § > Ad it Papel gra tao —0.54
ae VSG OS ie aa +2.7 || Br°+30H- .| 5 | BrO;-+3Ht |—0.60
oe i 2 |Catt+ +2.6 |) Fett. ..... 1 |Fet+t.....|—0.75
ee ea a" Po ola etapa y ot en | a. a See a Ot ie ‘1—0.80
MARE ee 3 | Alt++ 413 oHe® oF 2 |Hett+..... —~0.80
MPS es. 7 ed 9 a 4d dt BOs oo 2 | O.gas+2H* |—0.80
Peres oes. VA 07 ae ae +0.76'} Ch° gas+ 2 GHGS: 253 —0.85
Sa Ae 2 ae ba a ee +0.55 20H
: a Teale © aman lea 043i Het ii... a are grea —0.86
2 haa eee Pel oe ey +0.40}| 2Hgt....... 2 | 2Hgt++ —0.92
Pb°+SO0 = ed ig Pa +0.54)2Br-. oo... 7 Des 3.) ee ae ae ae —1.08
MUN ci ko 2° 1 Oor®.: +0.23)} Cl-+OH— BS DEON ss StS, —1.10
ees... Yay is, | a Saree +0.22/| I1°+30H- 5 |10;+8H*.. |—1.19
Cu°+20H— 2 | Cu(OH), +0.21|} Cré+++ HCrO;-+ {-1.3
A a Page i sel +0.12 4H.O| 3 4H+
Meee iss. . SRE ais. +0.10)} Mnt* + MnO:+ —1.35
a 3 | Fettt +0.04 2H20}| 2 4H+
eee ts.’ DS sieeet atiet +0.00)} 2CI-....... 7 Soe Nh ) PR ee —1.35
Ct. eek. 1 | Cutt... .}—0.17]| Pht ++ 2 | PbO.+2H+ |—1.44
Ae°+Cl- ...| 1 | AgCl....... ~ 0.23 20H -
Hg.Cl, soln. | 2 !2HgCh..... —0.24)) Au®......... j eS sae tes —1.5
+2Cl— PENS «et Loh, + Vad 61 (2 a ee —1.51
I°+OH— Oe is eee —0.26)} Mnt++ 5 | MnO, + —1.52
‘Os Tae 3 | Ast++ —0.29 4H.0 8Ht
ROU uk. ths > eg 1B ai apa ee —0.34!| MnO.+2H.0} 3 | MnO; + —1.63
Ag®+2NH;3 1 | Ag(NH;).+ .;—0.38 4H+
BERS es AES dar f SOON 2HeO .. 56 oh 2 | H.0.+2Ht .|—1.0
A eng alas 3 |Cot++. ....|—0.40]| Cott. ..... 1 |Cottt,...|/-1.8
0) 5 Meas 2 | O.+H,0 —0.41|| Cl, gas+ 4 | 2ClO~ + —1.8
Sane 3 |Sbtt++t+.....|—0.47 20H— 2Ht
Br®. liq.+ 5 | BrO;- +H,0|—0.51]| 30,2 gas...... 3 | 203 gas..... —1.9
30H— See a 2 Fe Ree —1.96

The table of oxidation potentials will help to explain many of the
reactions used in analytical chemistry. All the metals above hydrogen
will replace the hydrogen of dilute acids; those below hydrogen will
not do soasarule.. The oxidation of copper to cuprous ions, however,
corresponds to an oxidation potential of —0.17 volt. When the acid
is very concentrated and the cuprous solution very dilute, the oxida-
tion potential of the copper to univalent copper will be above that of
hydrogen and consequently a little copper should go into solution as
cuprous salt. In the presence of some oxidizing agent lower than

44 GENERAL PRINCIPLES

hydrogen ions in the series, the copper will dissolve readily. Such an
oxidization agent is the ferric ion, for the table shows that the oxida-
tion potential of ferrous ions to ferric ions is —0.75 volt; just as cupric
ions will oxidize metallic iron to the ferrous condition so will ferric
chloride oxidize metallic copper to the cupric condition:

Fett *+-+Cu > Cut t+2Fet tf

The table shows where interference is to be expected. Metallic
aluminium will precipitate iron from a solution of a ferrous salt. It
will not precipitate iron if an acid is present because, as the table shows,
iron will itself liberate hydrogen gas from hydrogen ions; on the other
hand, from the oxidation potential of ferrous to ferric ions, it is evi-
dent that metallic aluminium will reduce a ferric solution before
acting upon the hydrogen ions of the acid. It also shows that metallic
aluminium will precipitate copper completely even in the presence
of acid.

Lead and tin occupy neighboring positions in the series. In a
neutral solution, therefore, lead will precipitate tin from a solution
containing stannous ions. As the concentration of the lead ions in-
creases and the concentration of the stannous ions decreases, the oxida-
tion potentials approach one another so that equilibrium is soon reached.
Conversely, when no stannous ions are present, metallic tin will pre-
cipitate a little lead from a solution containing lead ions, but the equilib-
rium will soon be reached. ‘The presence of acid, however, will stop —
both of these reactions. On the other hand, the oxidation potential
of stannous to stannic tin, although not given in the table, has a nega-
tive value, and the lead will’reduce stannic ions even in the presence of
acid.

All reactions of oxidation and reduction represent reversible reac-:
tions. The strong reducing agents on being oxidized become weak
oxidizing agents, and conversely the strong oxidizing agents on being
reduced become weak reducing agents. Ordinarily hydrogen is
considered a reducing agent, but when a metal replaces the hydro-
gen of an acid, hydrogen ions act as the oxidizing agent.

The mass action law holds for oxidation and reductions as for all
other chemical reactions that take place in solution; when the oxida-
tion voltages are far apart, however, the reaction of oxidation and
reduction will apparently go to completion.

We shall now turn aside from reactions of oxidation and reduction
and consider some further applications of the mass-action law as applied
to reactions that are of importance in analytical chemistry.

THE EFFECT OF A COMMON ION 45

The Effect of a Common Ion

In all equilibrium expressions in which the concentration of an
ion is involved, the source of the ion is a matter of indifference. Thus
when sodium chloride and potassium chloride are both present in solu-
tion, the chlorine ion is common to each. The equilibria between

the salts and the ions are expressed as follows: 7
[Nat]x<[Cl]_, [K*] x[CI-] _
[NaCl] — tact and [KCl] = Mon

In each of these equilibrium expressions the concentration of the
chlorine is the total concentration in the solution. It is evident,
therefore, that the addition of the extra chlorine ions from potassium
chloride tends to lessen somewhat the extent of the dissociation of
the sodium chloride and similarly the presence of the chlorine ions from
the sodium chloride tends to lessen somewhat the extent of dissociation
of the potassium chloride. Both sodium and potassium chloride belong
to the class of strong electrolytes and the dissociation remains con-
siderable even after the other salt is added. As far as the chemical
behavior of the salts goes, the effect is scarcely noticeable. As long as the
solution is reasonably dilute the dissociation of both potassium and so-
dium chloride will be considerable in spite of the presence of the other salt.

The relations are quite different when one of the original substances
is difficultly soluble or only slightly ionized. Thus when a precipitate
of silver chloride is formed, the solution is saturated with the salt and
- the solubility product of the ions has been reached. If now a small
amount of either chlorine or silver ions is added to the saturated solu-
tion of silver chloride, further precipitation of silver chloride should
take place. As a general rule, therefore, a slight excess of precipitant
will make a precipitate less soluble. There are exceptions to this
rule, however. ‘The rule does not hold if the one of the ions from the
precipitate shows a tendency to form a soluble complex ion with the
excess of precipitant. Thus silver chloride is, in fact, less soluble in
very dilute sodium chloride solution than it is in water, but it dissolves
in a saturated brine solution more than in pure water, probably owing
to the formation of a complex ion. Barium sulfate is less soluble in
dilute sulfuric acid than it is in water, but concentrated sulfuric acid
dissolves it quite readily. Aluminium hydroxide is precipitated by
the very careful neutralization of an aluminium salt, but the hydroxide
dissolves in sodium hydroxide solution, forming sodium aluminate.
None of these examples is contrary to the mass action law, but they
show the need of considering all the possible reactions.

When a weak electrolyte is present in solution, the effect of the

46 GENERAL PRINCIPLES

common ion is often quite remarkable. To illustrate this effect let
us consider the weak electrolytes acetic acid, hydrogen sulphide and
ammonium hydroxide. The table on page 10 shows that a 0.1 N
solution of acetic acid is dissociated to between 1 and 2 per cent. The -
value of the ionization constant at 18° is about 0.000018. From this
value, the ionization (x) in 0.1 N solution can be computed as follows:

> =0.000018 - 2=0.00134 = 1.34 per cent.

The table on page 10 shows that the salts of the type represented by
sodium acetate are dissociated to about 84 per cent. A liter of 0.1 N
sodium acetate solution contains, therefore, 0.084 mole of acetate ions,
whereas one of 0.1 N acetic acid contains only 0.0013 mole of acetate
ions. When enough solid sodium acetate is added to 0.1 N acetic
acid to make the solution 0.1 N with respect to both the acid and
the salt, the common acetate ion tends to repress the ionization of
both the original molecules; the total concentration of acetate ion
is increased only very slightly and we are justified in assuming that
0.084 represents with sufficient accuracy the concentration of the
acetate ion. The concentration of the hydrogen ions may be des-
ignated again as x and that of the non-ionized acetic acid as 0.1—z.
The ionization constant of acetic acid remains the same and the
mass action expression becomes

0.084x
0.1—z

By adding an equivalent weight of sodium acetate, therefore, the ioniza-
tion of 0.1 N acetic a¢id is changed from 1.8 per cent to 0.02 per cent.

The effect of ammonium salt upon the ionization of ammonium
hydroxide is similar. The table on page 22 gives the solubility
product of ferric hydroxide as 1.110-%6 and that of magnesium
hydroxide as 3.410". As the cube of the OH” concentration is
taken in computing the solubility product of ferric hydroxide and
only the square of this concentration in the case of magnesium
hydroxide, the difference in solubilities is not as great as these
solubility products would indicate, but the values show that it takes
only an extremely low concentration of hydroxyl ions to satisfy the
solubility product of ferric hydroxide and many times as much to
satisfy that of magnesium hydroxide. Ammonium hydroxide added
to asolution containing ferric and magnesium ions will cause the pre-
cipitation of both iron and magnesium as hydroxide, but by adding
sufficient ammonium chloride to the solution, the ionization of the
base is repressed, exactly in the same way that sodium acetate acted

= (0.000018; x=0.000021 =0.02 per cent.

THE EFFECT OF A COMMON ION 47

upon acetic acid, to such an extent that none of the magnesium
is precipitated, although the precipitation of the iron remains
practically complete.

The effect of hydrogen ions upon the ionization of hydrogen sulfide
is similar. The solubility product of copper sulfide is 8.510-*5
and that of zine sulfide is 1.2<10-°%. The precipitation of both cop-
per and zinc by hydrogen sulfide is practically complete in a solution
containing no excess hydrogen ions. As the metal precipitates, how-
ever, hydrogen ions are formed:

Zn+H.S — ZnS+2H".

The accumulation of these hydrogen ions serves to repress the ioniza-
tion of hydrogen sulfide and tends to stop the precipitation of the zine.
If, therefore, we wish to precipitate copper and leave zinc in solution,
all that is necessary is to add a little acid at the start; in 0.3 N
_ acid solution the precipitation of the copper as sulfide is practically
~ eomplete, while little if any zinc sulfide is precipitated. On the other
hand, if some sodium acetate is added to the solution, non-ionized
acetic acid is formed and the accumulation of the hydrogen ions is
prevented. This effect is so remarkable that it is easier to precipi-
tate zine sulfide from a solution containing acetic acid and sodium
acetate than from a solution of zine chloride in water. The concen-
tration of hydrogen ions is kept very low, even although the solution
may smell strongly of non-ionized acetic acid.

The common ion effect is also involved in the solution of precipi-
tates. The effect is shown, for example, in the solution of calcium
phosphate. The table on page 10 shows that the tertiary ionization
of phosphoric acid is about the same as the secondary ionization of
hydrogen sulfide and the secondary ionization is comparable to the
primary ionization of hydrogen sulfide. If considerable hydrogen
ion is added in the form of a mineral acid, the ionization of HPO
and of HePO4 becomes extremely small. The saturated solution of
calcium phosphate contains Catt and PO.= ions. These PO4> ions
must be in equilibrium with added hydrogen ions. When the POs~
from HPO,” is kept less than the PO4~ concentration corresponding
to the value of a saturated solution of calcium phosphate, the calcium
’ phosphate must tend to dissolve. As Ca(HzPO.)2 is much more sol-
uble than CaHPOu, which in turn is more soluble than Ca3(PO4)g, it is
desirable to add enough acid to make the H2POx practically non-ionized.

The common ion effect which is involved in nearly every reaction
of precipitation and of solution represents an application of the mass-
action principle.

48 GENERAL PRINCIPLES

Hydrolysis

Hydrolysis is the name given to the decomposing action of water
upon many salts. Corresponding to the fact that water is a poor
conductor of electricity, it follows that water is dissociated only to a
slight extent.

H20 = Ht+0OH-.

According to Kohlrausch and Heidweiler,* the degree of disso-
ciation at 25° C. is 1.05107; in other words, about 10,000,000 liters
of water would furnish 1 gm. molecule of ionized water or 1 gm. atom
of hydrogen ions. Small as this is, it suffices to explain the hydrol-
ysis of:

I. The salts of weak acids with strong bases.
II. The salts of strong acids with weak bases.

III. The salts of weak acids with weak bases.

The salts of strong acids with strong bases are not hydrolyzed
appreciably. 3

Hydrolysis is shown to take place by the fact that solutions of
neutral salts corresponding to I react alkaline, those of II react acid, .
while those of III are sometimes acid and sometimes alkaline.

The cause of hydrolysis is the action of the ions of water upon
the ions of the dissolved salt.

All monobasic salts in aqueous solution are largely dissociated
into ions (cf. p. 10):

RA—Rt+A-,

and the phenomenon of hydrolysis may be represented by the general
equation
R*t+A~+H20 s$ ROH+HA.

ete — +’
Salt Base Acid

In formulating mass-action expressions involving highly ionized
and slightly ionized substances it has been found best to employ the
concentration of the ions of highly ionized substances, and the con-
centrations of the non-donized portions of slightly ionized substances.
Since in dilute solutions the mass of the water changes inappreciably
when the reaction takes place, we may neglect it in the formulation.

There are now three typical’ equations which represent the appli-
cation of the mass-action principle to the hydrolysis of the above
three classes of salts.

* Z. phys. chem., 14, 317.

HYDROLYSIS 49

I. Hydrolysis of Salis of Weak Acids and Strong Bases

The base formed by the hydrolysis is largely ionized, while the
acid is only slightly so. The mass-action equation now takes the
form |
[R*][OH [HA] _[OH7)[HA]

[R*][A7] 0s aa

Owing to the presence of the OH ions in appreciable amount,
all salts of this category react alkaline. The alkali salts of hydro-
eyanic acid, hypochlorous acid, carbonic acid, boric acid, and hydro-
gen sulfide are of this type.

dd) Ki=

II. Hydrolysis of Salts of Strong Acids and Weak Bases

Here the conditions are reversed, and it is the acid which is almost
completely dissociated and the base but slightly. In this case the for-
mula becomes

[ROH][H* [A] _[ROH)[A*]

[R*|[A7] [R*] ©

Salts of this class, such as those of chromium, aluminium, iron, etc.,
react acid when in aqueous solution.

In the case of polyvalent bases both electrolytic and hydrolytic
dissociation take place in stages:

(2) By=

RCls @ RClot+Cl-
meme | nog" eenci+c
| RCIt+2 Rttt+cr.

RClyt +H.0 @ RClLOH+ Ht
Hydrolysis RCI* ++2H20 = RCI(OH)2+2Ht
R**+*+3H20 = R(OH)3+3H".

Ferric, aluminium and chromic chlorides, for example, react acid
in aqueous solution. If such solutions be evaporated to dryness,
considerable hydrochloric acid is volatilized, and the residue obtained
in an insoluble Haste salt which can only be dissolved by means of
acid.

III. Hydrolysis of Salts of Weak Acids and Weak Bases

The acids and bases formed by hydrolysis are only slightly dis-
sociated, but to different extents. Our formula becomes

[ROH][HA]
[R*][A7]

(3) Ky=

50 GENERAL PRINCIPLES

Salts of this type are especially subject to hydrolysis and, as both
the acid and base are slightly ionized, the hydrolysis may take place
to a considerable extent without the soluten manifesting either acid
or basic properties.

If the electrolytic dissociation of the acid is greater than that
of the base, the solution reacts acid; and conversely, when the base
is stronger than the acid the solution of the salt shows an alkaline
reaction.

Neutral ferric acetate in a boiling, aqueous solution is hydro-
lyzed:

Fe(C2H302)3+2H20 @ Fe(OH)2C2eH302+2HC2Hs02.

Basic ferric acetate is precipitated and can be removed by filtering
the hot solution. If the solution is allowed to cool, the reaction tends
to take place in the reverse direction and some of the basic salt goes
into solution. Heat and dilution always favor hydrolysis.*

IV. Hydrolysis of Salts of Strong Acids with Strong Bases

Salts of this type yield, by hydrolysis, acids and bases which are
almost entirely dissociated in dilute aqueous solution, and the general
equation becomes:

[R* OH |[H*[A*]
[R*][Aq]

The H and OH ions, however, are in equilibrium with undissociated
water; the solution reacts neutral and contains only as many H and
OH ions as correspond to the ionization of water, which is so small
that there remains only the electrolytic dissociation of the salt to —
be considered.

Ky = =[H*][OH"].

Salts of this type are not subject to appreciable hydrolysis.

The hydrolytic action of water, as well as the mass-action law,
may be illustrated by the following experiment: A little water, added
to a solution of antimony chloride in concentrated hydrochloric acid,
causes precipitation of antimony oxy-chloride:

—Cl
Sb—Cl+}j, 02 = Sb_ oy +2HCI,

—Cl

*The ionization constant for water is 0.0:;12 at 25° but rises to 0.015 at 100°.
The water is, therefore, much more dissociated at the higher temperature and
as a result the hydrolysis is favored.

HYDROLYSIS 51

which dissolves on adding a little concentrated hydrochloric acid.

Further addition of water again precipitates the basic salt, which will
dissolve in more of the concentrated acid, etc. It is obvious that by
increasing the mass action of the water the reaction goes from left to

_ right, while by increasing the concentration of the hydrochloric acid
it goes from right to left.

The analytical chemist frequently desires to assist or to prevent

hydrolysis. The mass-action principle shows how this can be done.

To assist a chemical reaction it is necessary to increase the concen-
tration of one of the original reacting substances or to lessen the con-
centration of one of the substances formed. Aside from the reac-
tions of oxidation and reduction, all the reactions that take place
completely in aqueous solution are those in which the concentration of
one of the substances formed is practically negligible. This is due to
the formation of (1) a precipitate, (2) a gas, or (3) an undissociated sub-
stance (cf. p. 26). Neutralization takes place betweenanacid and a base
because of the tendency to form undissociated water. Hydrolysis is
the reverse of neutralization and is due to the fact that water is slightly
dissociated. Hydrolysis takes place when one of the products is a
gas, a precipitate, or an undissociated substance. Thus the hydrolysis
of a salt of a weak acid and a strong base is due to the fact that the
dissociation of the weak acid is slight. The hydrolysis of a salt of
a strong acid and a weak base is due to the formation of the undisso-
ciated base. Hydrolysis takes place most readily when both the acid
and the base are weak, because then both the H ions and the OH ions
are removed from the solution to form undissociated acid and undisso-
ciated base. Hydrolysis of a salt of a strong acid and a strong base
cannot take place because there is then no tendency for the H and OH
ions of water to be removed. ‘To assist hydrolysis, boiling is advisable,

- because the water is so much more dissociated at this temperature

than when cold. The reaction that takes place on boiling often
proceeds in the other direction on cooling, simply on account of the
change in the ionization of the water itself. Dilution favors hydrolysis
because it diminishes the concentration of the substances formed; the
concentration of the ions from water is not changed, but the relative pro-
portions of these ions to other ions present is increased. Hydrolysis is
also favored when one of the products is gaseous or insoluble; thus
ammonium carbonate is very easily hydrolyzed, because both ammonia
and carbon dioxide are volatile, and ferric acetate is hydrolyzed not
only because the acid and base are both weak, but the basic ferric
acetate is also very insoluble. Finally, to stop hydrolysis it is only
necessary to add a little free acid or free base at the start, and the

52 GENERAL PRINCIPLES

efficiency of the acid or base is proportional to the extent to which
it is itself ionized. If, the solution becomes alkaline as a result
of hydrolysis, then a little alkali will best stop hydrolysis, but if the
solution becomes acid, a little acid should be added. Moreover, it is
advisable to work in cold and concentrated solutions. A dilute ©
aqueous solution of potassium cyanide has the odor of free hydro-
cyanic acid and reacts alkaline to litmus, but if a little caustic potash —
is added, the hydrolysis of the salt is prevented and the odor of hydro-
cyanic acid can be detected no longer. Similarly, boiling a solution
of neutral ferric sulfate results in the precipitation of a basic salt, but
the precipitation does not take place if a little sulfuric acid is added
at the start. |

Amphoteric Electrolytes

According to ideas that prevailed nearly a century ago, an acid was
regarded as composed of a negative element and hydroxyl; a base was
regarded as composed of a positive element and hydroxyl. According
to the modern conception, an acid is a substance capable of yielding
hydrogen ions and a base is a substance capable of yielding hydroxyl
ions. The so-called negative elements are sometimes simple negatively-
charged ions, as the Cl” of hydrochloric acid, and sometimes complex
anions, such as the SO," of sulfuric acid, containing a positively charged
element and negatively charged oxygen.

The extent to which the ionization takes place in the solution |
of an acid or a base varies greatly. In 0.1 N solution, the table
on page 10 shows that some acids and some bases are dissociated to
about 90 per cent, while other acids and other bases are dissociated
to 1 per cent or less. We are accustomed to regard the strength of
an acid or of a base as shown by the extent to which ionization takes
place in aqueous solution.

As a general rule, when the positive charge on an element is in-
creased, the more difficult it becomes for the corresponding hydroxide
to ionize as a base and the greater the tendency for the hydroxide ©
to ionize as an acid. Thus the higher the positive charge on the
atom, the more negative it becomes according to the old-fashioned
conception. This apparent contradiction is not quite as inconsistent
as it seems; it arises from a confusion of the intensity factor of
electric energy with the quantity factor. All forms of energy are
composed of these two factors and such confusion of the factors is quite
common.

The quantity of electricity corresponding to a unit electric charge
on a gram atom of any element is 96,500 coulombs. All univalent

AMPHOTERIC ELECTROLYTES Ne BS

a elements bear this charge and all bivalent elements twice as much.

_ On the other hand, the tendency of the atoms to accept or give up

electrons varies greatly, as the table of oxidation potentials shows
(p. 48). It is customary to speak of the elements which show the

| greater tendency to give up one or more electrons of negative electricity
| as more positive than those which show less tendency to lose electrons.
_ In other words, the elements nearer the top of the electromotive series
are regarded as more positive than those which are below them in this

series. ‘Thus the older conception of positive and negative nature
was largely one of electromotive force, or voltage, rather than one of
electric nature or quantity of electricity. It corresponds to the
behavior of the element in an electric couple such as that of the Daniell
cell; the zinc becomes the anode in such a cell and the copper the
cathode; the current flows through the cell from the zine to the cop-
per and it is customary to say that the zinc is positive to the copper.
The student should be careful to distinguish in his own mind the dif-
ference between the meaning of the word positive when used in this
sense of voltage and when used to represent the character of an elec-
tric charge.

Water is a substance which ionizes to a slight extent and the initial
products of the ionization are H+ and OH-. Water, therefore, may
be considered as being both an acid and a base. There are other
substances which act as acids without forming at one time any more
hydrogen ions than does water and other basic substances which do
not form more hydroxyl ions. Water is characterized by the fact
that it forms an equal amount of both hydrogen and hydroxyl ions
by its primary ionization. Other hydroxides are known which have
both acid and basic properties, but this is due to two distinct kinds of
ionization. At one time they dissociate as an acid and at another
time as a base. Such substances are said to be amphoteric electrolytes.

Aluminium hydroxide is an amphoteric electrolyte. This sub-
stance has a very small solubility product and the quantity of ions
present in the saturated solution’ is extremely small. The chemical
behavior of aluminium hydroxide shows; however, that it is capable
of dissociating in two ways:

Al(OH)3 — Al(OH)2*-+OH™ (ionization as a base);
Al(OH)3 — Ht+AlO3H2- — Ht+AlO2~+He20 (ionization as an acid).

The mass action principle enables one to predict which of these
ionization reactions will take place. In the presence of a strong acid,
such as hydrochloric acid, the common ion effect of the hydrogen

54 GENERAL PRINCIPLES

ions will prevent the ionization of the aluminium hydroxide as an
acid; the effect is much more marked than that of sodium acetate
upon acetic acid, for in this case the difference in the percentage
ionization of hydrochloric acid and aluminic acid is much greater.
On the other hand, the presence of the hydrogen ions favors the ion- ~
ization of the aluminium hydroxide as a base. To establish the proper
equilibrium between Ht, OH~ and H20, nearly all the OH™ ions from
the aluminium hydroxide react with the H* of the hydrochloric acid.
On account of this removal of the OH~, the ionization of the Al(OH)s
progresses, the Al(OH)2* ionizes into Al(OH)t* and OH™ and finally
the Al(OH)** ionizes into Al**+ and OH. In this way Al(OH)s3
dissolves to form AICls.

Similarly, in the presence of sodium hydroxide, the common ion
effect. of the OH™ prevents the aluminium hydroxide from ionizing
as a base and the H* ions formed by its dissociation as an acid react
with the OH™ ions from the sodium hydroxide, and the final result is
that the aluminium all dissolves as sodium aluminate, NaAIOzg.

Whenever an oxide or hydroxide dissolves by chemical reaction
with an acid and also by chemical reaction with a base, the substance
is obviously an amphoteric electrolyte.

Detection of Acids and Bases. Theory of Indicators *

To detect the presence of free hydrogen cations or of hydroxyl
anions, certain colored, organic substances are used, called indicators.
These indicators are very weak acids, very weak bases or amphoteric
electrolytes and the free acid, or base, is an unstable substance
which tends to undergo a slight rearrangement of the atoms in order
to assume a condition of greater stability. The color of any organic
compound is due to a certain special arrangement of certain atoms,
the chromophor; when this arrangement is changed, the color is also
changed or lost.

Methyl orange is an amphoter and is capable of forming salts
with both acids and bases, but its indicator characteristics are due
to its very weak basic properties. The neutral solution of its sodium
salt is used as an indicator. In this sensitive neutral solution we have
a condition of equilibrium between the two isomeric forms of methyl
orange as expressed by the equation

HSO3 E CgHsN :N- CeH4N(CHs)2 = OsteBane -N: CeH4 : NC

*Cf. J. Srizaciitz, J. Am. Chem. Soc., 25, 1117; 39; Acres, tbid., 37, 39, 42;
McCoy, ibid., 31, 508; Saum, Z. phys. Chem., 57, 471; and A. A. Noyzrs, J. Am.
Chem. Soc., 32, 815.

DETECTION OF ACIDS AND BASES 55

_ The compound on the left is yellow in color and its color is due to
the azo group N : N; the other compound is red, having for its chromo-
phor the quinoid group : CgHa:.
The sodium salt of methyl orange is yellow and has the formula
NaSO3 : CeHaN :N- CeH4N(CHs)z2,
and when decomposed by acids the free sulphonate
tril on gaa : CeH4 ; N(CHs3)2
|

is formed, which is red.

The red quinoid form is ionized as a weak base and forms red salts
with acids. It does not form salts readily with weak acids, such as
carbonic or acetic acid, because, as we have seen, salts of weak bases
and weak acids are hydrolyzed. This is why methyl orange is not a
sensitive indicator for weak acids. As a very weak base it will be driven
readily out of its red salts by other bases, even weak ones; and the free
base will revert again to its yellow form, the result being that methyl
orange is an excellent indicator for weak bases.

Phenolphthalein, another valuable indicator, is a very weak
acid. The free acid, however, is unstable, and when set free from
one of its colored salts reverts instantly into a colorless lactoid form:

HOOCCsHsC(CeH.OH) . Caki«:s Me See Cer as
Re

Colorless

The red color is in this case also due to the quinoid grouping: CgHa:.
In the free acid, the condition of equilibrium favors the lactoid form,
and only minimal quantities of the quinoid acid are present. This
trace of quinoid acid is ionized and in equilibrium with its ions:

HOOCC.H4C(CgH40H) : CeHs : OR
Ht+O0CC.gH1C(CgH:0H) > CeHs: OO.

The addition of an alkali causes the hydrogen ions to disappear,
more of the quinoid molecules must be ionized to preserve equilib-
rium, and the quinoid molecules in turn be reproduced from the lactoid
as fast as the former are converted into the salt. Phenolphthalein is
@ very sensitive indicator toward acid, but on account of being such
a weak acid it does not form stable salts with weak bases.

Besides these indicators, others are often employed, among which
may be mentioned Litmus and Lacmoid, which are red with acids
and blue with alkalies; and Turmeric, which is brown with alkalies
and yellow with acids.

‘

56 GENERAL PRINCIPLES

The various indicators, therefore, differ from one another with
regard to the extent to which they normally undergo ionization either
as an acid or as a base. They show, as a result, specific degrees of
sensitiveness to hydrogen cations and to hydroxyl anions. The follow-
ing table shows the concentration of the ion required to produce the
color change. The tables refer to results obtained when about two
drops of a 0.1 per cent solution of the indicator is added to 10 ce. of |
the solution tested.

SENSITIVENESS OF INDICATORS *

Concentration Required
for Color Change of
Indicator. Color with H+. Color with OH-.
Ht, OH-.
Phenolphthalein.....| colorless........ pink (3 PA, 0.071 0.041
Azolitmin (in litmus).| violet pink...... WIGOLS «hae s 0.051 0.051
Methyl orange...... reddish-orange..| yellow....... 0.0.1 0.01

The table shows that of these three indicators phenolphthalein is
the most sensitive to acids, and methyl orange is the most sensitive to
bases. |

The Evaporation of Acids

It is often necessary to diminish the acid concentration of a solu-
tion. The simplest way to accomplish this is by neutralization, but
it is often undesirable to introduce foreign substance into the solution,
and in such cases the acid is removed by evaporation. To prevent loss
by spattering, it is well to evaporate in a porcelain dish on the steam
bath; but, to save time, the evaporation may take place over a free
flame. In this case the solution should be kept in motion either by
stirring or by rotating the contents of the dish and moving it back and
forth over the flame. For evaporations over a free flame, an Erlen-
meyer flask is often used; the’sides of the flask, being cooled by the
air, act as a condenser and delay evaporation, but they also serve to
prevent mechanical loss. Spattering is caused by steam being formed
too rapidly at the bottom of the dish, by steam being formed beneath
a crust which has formed on the surface of the liquid, or by steam being
formed from enclosed mother-liquid in a solid that has separated. The
steam escapes then only after its pressure has become considerable.
The breaking of glass dishes when heated over a free flame is due to
overheating the glass by contact of the flame with some part that is

*From J. Stieglitz: Qualitative Analysis.

THE EVAPORATION OF ACIDS | 57

not being kept relatively cool by contact with liquid. Moving the
vessel back and forth over the flame prevents breakage by preventing
overheating of the glass; the bottom is kept uniformly wet on the inside
even when but little liquid remains.

The behavior of acids upon evaporation is interesting. A solution
of hydrochloric acid of specific gravity 1.10-boils at 110° under atmos-
pheric pressure. The solution contains 20.2 per cent of hydrogen
chloride and its concentration is thus a little less than 6 N. No
other mixture of water and hydrochloric acid boils as high and con-
sequently in evaporating a solution of hydrochloric acid, although the
total quantity of hydrochloric acid constantly diminishes, the concen-
tration of the remaining acid will tend to approach the composition of
the constant boiling mixture. If the acid is more dilute at the start,
'_ evaporation will cause the acid that remains to be more concentrated,
and if more concentrated at the start, evaporation will cause the

remaining acid to be more dilute. .

Nitric acid similarly forms a constant boiling mixture with water.
This mixture boils at 120.5° and contains 68 per cent of nitric acid;
the acid concentration of the mixture is about 15.5 N.

Sulfuric acid forms with water a constant boiling mixture that con-
tains 98.3 per cent of anhydrous sulfuric acid and boils at 330°. It was
once thought that these constant boiling mixtures of acid and water
corresponded to definite hydrates of the acid, but the composition of

such a mixture and the boiling-point varies with changes in pressure in a
way that proves no definite compound is present.

When it is desired to change a solution of a chloride to one of a
nitrate it is very easy to do this by evaporating once or twice with
nitric acid. ‘The nitric acid reacts with the hydrochloric acid,

6HCI+2HNO3 =4H20+2N0+3Clh,

* » and by adding an excess of the nitric acid all of the chloride is decom-
: posed. The most economical way to accomplish the change is to
evaporate nearly to dryness and then add a little strong nitric acid,
repeating the evaporation and addition of acid until no more red fumes

are evolved on adding the acid (NO-+air-red NOz).

Similarly a solution of nitric acid may be changed to one of hydro-
chloric acid. Aqua regia is formed, as in the above case, and by using
an excess of hydrochloric acid all the nitrate is decomposed.

The change of a solution of nitric acid or of hydrochloric acid to
one of sulfuric acid is based on another principle. In this case the

change is accomplished by evaporating till dense fumes of sulfuric acid

58 GENERAL PRINCIPLES

are evolved. Very little sulfuric acid is lost until all of the more volatile
acid has evaporated. |

Phosphoric acid is still less volatile. By evaporation with phos-
phoric acid even sulfuric acid can be distilled off without losing an
appreciable quantity of phosphoric acid, but the phosphoric acid is
changed to pyrophosphoric acid or to metaphosphoric acid.

To change a sulfate or a phosphate solution to a chloride or nitrate
presents a more difficult problem than that of simple evaporation. It
is customary in such cases to remove the phosphate or sulfate ions by
precipitation or to precipitate the desired substance and to dissolve it
in the desired acid.

Filtration and Washing of Precipitates

When a precipitate is produced in qualitative analysis it is necessary
to remove it from the liquid by means of filtration. The size of the filter
used should be determined by the size of precipitate and not by the
volume of liquid. In attempting to detect the presence of traces of a
substance it is often necessary to work with large quantities of the orig-
inal substance, and this involves the use of correspondingly large
volumes of liquid. If from such a solution a few milligrams of a pre-
cipitate is formed, it will be practically lost if spread over a large filter,
and further work with the precipitate is hampered.

Before examining such a precipitate it must be completely freed
from all traces of the filtrate. This is accomplished by washing. Wash-
ing must be continued until no test can be obtained with the wash-
water for a certain substance known to be present in the filtrate. For
example, suppose it is necessary to filter off some suspended barium sul-
fate from a solution containing sodium sulfate; the precipitate must be
continuously washed until a sample of the wash-water, acidified with
hydrochloric acid, no longer gives a precipitate on the addition of barium
chloride. As a rule, it is not advisable to run the wash-water into the
filtrate, because this occasions an unnecessary dilution of the latter.
The filter must always be smaller than the funnel, and the precipitate
should not extend higher than to within 5 mm. of the top of the filter.
Finally, large precipitates should be avoided as far as possible, for they
render exact work more difficult—filtration and washing consuming
too much time.

Colloidal Solutions

If an alkaline silicate in dilute solution is mixed with dilute hydro-
chloric acid, the solution remains clear and apparently unchanged in

hE a a

COLLOIDAL SOLUTIONS 59

spite of the fact that silicic acid is only slightly soluble in water. If the
liquid thus obtained is placed in a vessel whose walls are formed of
parchment paper or of bladder and this vessel is placed in pure water,
it is possible to wash out all the excess hydrochloric acid and all of the
sodium in the form of sodium chloride. This is the so-called process
of dialysis. The liquid in the dialyzer contains silicic acid.

‘The silicic acid, however, is not in the form of a true solution. If it
is evaporated, an amorphous, gelatinous mass is obtained which is only
slightly soluble in water. The boiling-point and freezing-point of the
solution do not differ much from that of pure water. The addition of
various substances, especially of salts, causes the liquid to solidify as a
jelly, especially if it has been concentrated somewhat by evaporation
in the cold.

A similar experiment can be performed with arsenious acid and
hydrogen sulfide. The aqueous solution of arsenious oxide turns
orange yellow when hydrogen sulfide is passed into it and becomes
opalescent without forming any precipitate, although arsenious sulfide

is only very slightly soluble in water. The liquid will pass unchanged

through an ordinary paper filter. If a little hydrochloric acid or some
salt solution is added to the yellow opalescent liquid, a heavy precipi-
tate of arsenious sulfide at once forms.

A liquid in which a very insoluble substance appears to be in solu-
tion far beyond its usual degree of solubility, and yet does not show at ~
all the behavior of an ordinary supersaturated solution, is said to con-
tain the substance in colloidal solution.

Such solutions are very commonly formed with difficultly soluble
gelatinous substances. It was formerly thought that such solutions
represented true solutions of difficultly soluble substances in the form
of a soluble colloidal modification, but more recent observations with the
ultramicroscope have indicated that colloidal solutions are really sus-
pensions of minute solid particles. The colloidal condition evidently
interferes with the precipitation of insoluble substances, and since most
of the separations of analytical chemistry depend on the formation of
precipitates, analytical chemistry is chiefly concerned with the colloidal
condition as one that is to be avoided as much as possible.

The suspended particles of most colloids carry electrical charges;
a potential difference exists between the particles and the liquid. Some-
times the charge on the particles is positive, as with the colloidal solution
of ferric hydroxide, aluminium hydroxide and chromium hydroxide;
and sometimes the charge is negative, as with the colloidal solu-
tions of silicic acid, stannic acid, the sulfides of arsenic and cadmium,
silver iodide, silver chloride, gold, platinum and silver. Substances

60 GENERAL PRINCIPLES

in the colloidal condition which carry an electric charge are precipitated
by an electrolyte. Negatively-charged colloids are precipitated by the
action of positive ions and, conversely, the positively-charged colloids
are precipitated by negative ions. The precipitated substance carries
with it a little of the precipitating ion in the form of an adsorption com-—
pound. The precipitating power of electrolytes increases decidedly
with the valence of the precipitating ion.

Some colloids do not carry much of an electric charge and sometimes
the charge is positive in acid solutions and negative in alkalies. Col-
loids which do not carry much electric charge are not precipitated
by dilute solutions of electrolytes. These colloids are coagulated by
heat, by any other method of dehydration such as the addition of con-—
centrated salt solutions, or by the addition of another solvent such as
alcohol. Colloids which are not sensitive to precipitation by electro-
lytes often act as protective colloids and hinder the precipitation of
other substances. Thus tannic acid, gelatin and albumin act as such
protecting agents, probably by forming protective films.

To prevent the colloidal condition in analytical work it is necessary
to take certain precautions. *Since electrolytes cause the precipitation
of colloids, it may be expected that their absence will tend to aid in
the formation of colloidal solutions. Thus in washing a precipitate
which is known to have a tendency to pass into colloidal solution, it is
best to wash with a solution of some indifferent electrolyte rather than
with pure water. Again, if precipitations are attempted in very dilute
solutions of such substances as arsenious sulfide, it is best to have an
electrolyte in the solution at the start. It is also necessary to remember
that the colloids carry down with them the precipitating ion by which
they are coagulated. To avoid analytical mistakes caused by such loss
of ions, a sufficient concentration of some other ion is usually provided
in the form of an acid or an ammonium salt. The washing of the
precipitated colloid with ammonium chloride or ammonium nitrate
gradually removes these precipitated ions, but it is very difficult to avoid
all danger of loss from this source.

When protective colloids, especially of the gelatin or albumin type
are present, they may interfere so much with the common precipitation
tests that it is necessary to destroy the protective colloid, usually by
oxidation, before proceeding with the analysis.

.
v

REACTIONS IN THE DRY WAY 61

II. REACTIONS IN THE DRY WAY

These reactions are employed chiefly in the so-called ‘‘ preliminary
examination,” in testing the purity of precipitates, and in the examina-
tion of minerals. The most important reactions of this nature consist
in the testing of a substance with regard to its—

1. Fusibility;

2. Ability to color the non-luminous Bunsen flame;
3. Volatility;

4. Behavior toward oxidation and reduction.

In order to carry out these reactions it is customary to use the
non-luminous gas flame; and to understand the operations to be
described it is necessary for us to know something about the composi-
tion of illuminating gas and the nature of the flame.

The illuminating gas of Zurich averages the following composition:
CO2=2.0 per cent; C,H2,=4.5 per cent; O2=0.2 per cent; CO=8.0
per cent; H2=48.0 per cent; CH4=33.0 per cent; and No=4.3 per
cent.

All these components, except CO2, O2, and Ne (which are present
only in small amounts), are combustible; they are reducing substances.
illuminating gas ordinarily burns with a luminous flame, and the
luminosity is due to the presence of unsaturated hydrocarbons (C,Ho,),
principally ethylene, propylene, acetylene, benzene, etc. If ethylene
is heated to a certain temperature, it is decomposed into methane and
carbon:

C2oH4=CH4+C,

and it is glowing carbon which causes the luminosity of the flame. |

The other unsaturated hydrocarbons behave like ethylene. The
remaining combustible constituents of illuminating gas burn with a
non-luminous flame. If we bring air into the gas, the flame becomes
non-luminous. With the Bunsen burner air is introduced by opening
the holes at the base of the burner. In such a gas-flame there are,
according to Bunsen, the following parts (Fig. 1*):

I. The inner cone of the flame, aab, in which no combustion takes
place, because the temperature here is too low. This part of the flame
contains unburned gas mixed with about 62 per cent of air.

II. The flame mantle, indicated by acaba, which is composed of
burning gas and air.

*In the drawing, the lines d represent a metallic chimney, or flame-protector,
which rest upon the support ee. It is advisable to furnish each burner with such
a flame protector.

62 GENERAL PRINCIPLES

Ill. The luminous tip, at b, which does not appear unless the air-
holes are closed somewhat.

In these three principal parts of the flame Bunsen distinguished
Six reaction zones: |

1. The base of the flame at a’. The temperature here is relatively
low, because the burning gas is cooled by the constant current of
fresh air, and also because the burner itself conducts away considerable
heat. This part of the flame serves to test volatile substances to see
whether they impart color to the
flame. In case several substances are
present which color the flame, it is
often possible to observe the colors
one after the other, in that the most
volatile substance colors the flame
first, and later the colors caused by
the less volatile ones are seen. This
would not be possible at a hotter part
of the flame, as all of the substances
would then be immediately volatilized,
producing a mixture of colors.

2. The fusing zone at B. This lies
at a distance of somewhat more than
one-third of the height of the flame,
and equidistant from the outside
and the inside of the mantle, which
is broadest at this part. As this is
the hottest part of the flame (about
2300° C.), it serves for testing sub-
stances as to their fusibility and.
i volatility.

Fic. 1. 3. The lower oxidizing flame lies in

the outer border of the fusing zone at

y, and is especially suited for the oxidation of substances dissolved
in vitreous fluxes.

4. The upper oxidizing zone, at ¢, consists of the non-luminous
tip of the flame, and acts strongest when the air-holes of the lamp
are fully open. It is used for various oxidizing tests, the roasting
away of volatile products of oxidation, and generally for all processes
of oxidation where the very highest temperature is not required.

5. The lower reducing zone lies at 6, in the inner border of the fusing
zone next to the dark cone. As the reducing gases are mixed here with
oxygen from the air, many substances which are reduced by the upper

©)

REACTIONS IN THE DRY WAY 63

reducing flame are unaffected in this zone. This part of the flame is
consequently very well adapted for a test which cannot be made
with the blowpipe, namely, reduction on the charcoal stick, and in
vitreous fluxes.

6. The upper reducing flame is at ny, in the luminous tip of the dark
inner cone, which may be produced by gradually diminishing the supply
of air. If the luminous tip has been made too large, a test-tube or
porcelain dish filled with water and placed over it will be blackened,”
which should never be the case. This luminous tip contains no free
oxygen, is rich in separated incandescent carbon, and has, therefore,
a much stronger reducing action than the lower reducing zone. It is
used more particularly for the reduction of oxides collected in the form
of incrustations.

METHODS FOR THE EXAMINATION OF A SUBSTANCE IN THE Dry Way

1. TEestT OF THE FUSIBILITY

This test is principally made in the examination of minerals, which
are introduced into the flame in the loop of a platinum wire (about as
thick as a horsehair). The sample is examined, after heating, by means
of a magnifying glass to see whether the corners are rounded, due to
melting. The potentially hottest temperature of the fusing zone
amounts to about 2300° C.* It will never be possible to reach this
temperature with the test, because the substance itself loses heat by
radiation. As the amount of heat lost by radiation is proportional
to the surface exposed, it is evident that we will obtain the maximum
heat by using a very small sample and holder. For this reason a coarse
wire should not be used for this test. |

We distinguish the following degrees of heat:

1. Faint red glow.......... 525°C. Melting-point of magnesium.. 631°
2: Dark red glow:..:....... 700° aluminium 658.7°T
3. Bright red glow.......... 950° silver..... 960.5°
TRO: 1063°
Moe GOW PIOW So eee ds 1100° Soca. 1083°
5. Faint white glow........ 1300° nickel.... 1452°
6. Full white glow.......... 1500° platinum., 1755°

* This temperature will be considerably lower with too large a supply of air.
According to Naumann, the temperature of illuminating gas with 1} times its
volume of air reaches about 1818° C., but the temperature obtained is usually
lower owing to loss by radiation. The finest platinum wire can be melted by
means of the flame, but not when it is as thick as a horsehair.

+ Circular No. 35 of the Bureau of Standards, Washington, D.C. Cf. Burgess-
IeChatelier: “ High-Temperature Measurements.”

64 GENERAL PRINCIPLES

Below 525° C. the following substances melt: tin at 232°, bismuth —
at 271°, lead at 327°, zine at 420°.

2. CoLoR IMPARTED TO THE FLAME

The substance (best in the form of the chloride) is placed in the
loop of a fine platinum wire, introduced into the base of the flame,
and then finally brought into the fusing zone.

3. TEST OF THE VOLATILITY

A little of the substance is heated in a small test-tube or in a piece of
glass tubing sealed at one end. Volatile substances are transformed
into vapor, often without melting, and the vapors condense on the
cooler walls of the tube.

4. OXIDATION AND REDUCTIONS

(a) In Vitreous Fluxes or Beads

To make a bead, borax (Na2B407:10H2O) or salt of phosphorus
(NaNHs4HPO4,+4H20) is used. A piece of very fine platinum wire,
about 3 em. long, is sealed into the end of a glass tube. The wire is
heated to redness, and then quickly dipped into the borax or salt of
phosphorus, held near the flame, whereby a small amount of the
salt isfused to the end of the wire. By repeated heating and dip-
ping into the salt a bead of sufficient size is obtained. This should be
about 1.5 mm. in diameter at the most.. It is not advisable to make a
loop at the end of the wire, because in this way the exposed surface is
unnecessarily increased. There is no danger of the bead falling off,
provided the wire is held horizontally in the flame and the bead is not
too large. In order to bring the substance in question into the bead,
it is only necessary to moisten the latter with the tongue, and then dip
it into the finely powdered substance, which will cause a small amount
to adhere to the bead. It is preferable to introduce too little substance
into the bead rather than too much, because, in the latter case, the bead
will become dark and opaque. The oxidation of the substance in the
bead is brought about by heating it in the lower oxidizing flame; reduc-
tion is usually effected by heating in the lower reducing zone, and cool-
ing in the dark inner cone, in order to prevent oxidation, which might
take place if the substance were cooled in the air.

In order to clean the wire, a borax bead is produced on the wire,
which is then heated, as is shown in Fig. 2, a, on one side of the bead

- REACTIONS IN THE DRY WAY 65

only, so that the latter runs along the wire in the opposite direction,
dissolving off all impurities. By heating the bead from the other
side, Fig. 2, 6, it is driven toward the end of the wire, from which it
can be shaken off by a quick jerk. By repeating this process three
times the wire is cleaned with the exception of a small amount of
adhering borax-glass, which can be removed by heating the wire in
the fusing zone until the sodium flame entirely disappears.

Fig. 2.

(b) Reduction on the Charcoal Stick

These exceedingly beautiful reactions are among the most sensi-
tive of those used in analytical. chemistry, and should be faithfully
practiced by every beginner. The cause of their sensitiveness is '
due to their taking place on the extreme end of a tiny piece of char-
coal, that is at a point, so that the sample has no opportunity to spread
itself over a large surface, which is the case with the ordinary reactions
on charcoal before the blowpipe.

To carry out these reactions, we use an ordinary splinter consisting of
good, straight fibers, such as used in an old-fashioned brimstone match
(not a safety-match, which has already been subject to chemical treat-
ment). It isimpregnated with sodium carbonate (soda) in the following
manner: A crystal of sodium carbonate (NazCOs-10H2O) is warmed
in the flame, whereby it melts in a part of its water of crystallization.
Three-fourths of the length of the match is now smeared with this liquid
soda, and the match is then slowly rotated on its axis in the flame,
until the soda:melts and penetrates the charcoal. On withdrawal
- from the flame there should be no place which continues to glow;
should the latter be the case, the stick should be quickly immersed in

66 GENERAL PRINCIPLES

the soda again. In this way one obtains a solid little piece of charcoal,
which can be heated for a long time without burning through.

In order to carry out a reduction, a small amount of the substance
to be examined is mixed on the palm of the hand with an equal amount
of calcined soda, a small drop of melted soda is added, and the mixture
is made into a paste by means of the blade of a penknife. The warmed —
piece of charcoal is then rubbed into the mixture, which adheres to it:
The sample is first heated in the lower oxidizing flame until it has melted,
and then moved into the lower reducing flame. The reduction will
be made evident by a violent swelling up of the melt, caused by the
evolution of carbon dioxide. As soon as the mass melts quietly the
reduction is complete. The substance is allowed to cool in the dark
cone, after which it is removed from the flame. The metal is now found
on the extreme end of the carbonized match, concentrated in a point.
This point is broken off, and triturated with a small amount of water
in an agate mortar. The excess of sodium carbonate goes into solution,
part of the charcoal floats on the surface of the water, while the heavier
metal sinks to the bottom. In case the reduced metal is iron, nickel,
or cobalt, it will not be noticeable to the eye, but it may be taken up
with a magnetized knife-blade, to which it will adhere, usually mixed
with charcoal. This should be dried by cautious warming, the tuft
of metal taken off, rubbed between the thumb and forefinger, and
then brought into contact with the knife again, to which only the metal
will now adhere. The metal is then transferred to a piece of washed-
filter-paper about 3-4 mm. wide and 50 mm. long, so that it comes as
near as possible to the end of the strip. By means of a capillary tube,
a drop of hydrochloric acid and one of nitric acid are added, and the
paper is warmed over the flame until the black speck (the metal) has
disappeared, when the final test can be made.

In order to test for iron, a drop of potassium ferrocyanide is added,
whereby the presence of iron is shown by the appearance of a distinct
formation of Prussian blue. To test for nickel and cobalt, the metal is
dissolved in nitric acid, the excess of acid is evaporated off, and a drop
of concentrated hydrochloric acid added, whereby the paper is colored
blue if cobalt is present; the nickel shows at the most only a very weak
greenish color—usually, however, no color. A little caustic soda
solution is now added, and the paper held in the vapors of bromine;
in case either nickel or cobalt is present a brownish-black spot appears,
due to the formation of either Ni(OH)3 or Co(OH)s.

If, however, the metal reduced was malleable, it is usually obtained
in the form of a metallic globule on the end of the match, where it can
be examined with the aid of a lens. Copper is not always obtained as a

i
o
4
;

REACTIONS IN THE DRY WAY 67

globule, but usually as a reddish, sintered mass. By pressing down on
a malleable metal in the agate mortar it is obtained as a glistening
fragment, which can be readily separated from the specifically lighter
charcoal by washing. To accomplish this the agate mortar is inclined
and a stream of water is directed sideways upon the mass, whereby
the charcoal is washed out with the water, and the metal is left clean.
It is transferred to a watch-glass and tested as follows:

1. The Metal is White (Pb, Sn, Ag, Pt). The metal is treated
with a few drops of nitric acid and carefully warmed. Lead and silver
dissolve readily, particularly upon addition of a little water. Silver
will be detected by the addition of a drop of hydrochloric acid, whereby
white silver chloride, soluble in ammonia, is precipitated. The test
for lead is dilute sulfuric acid, which precipitates white lead sulfate.

If the metal, on treatment with nitric acid, remains unchanged,
it is probably platinum. It should be dissolved in aqua regia, evapor-
ated to dryness, dissolved in a little water, and potassium chloride solu-
tion added. A yellow, crystalline precipitate confirms the presence of
platinum. Ifthe metal, when treated with nitric acid, becomes changed
into a white insoluble oxide, it is tin. In this case, another fragment of
metal is dissolved in concentrated hydrochloric acid and tested for tin
by means of mercuric chloride solution, or by a solution of bismuth
oxide in caustic soda.

2. The Metal is Yellow to Red (Cu, Au). Copper is readily dissolved
in nitric acid, and the solution gives with potassium ferrocyanide a

reddish-brown precipitate. Gold is insoluble in nitric acid, but soluble

in aqua regia. The evaporated solution gives a violet-brown color with
stannous chloride, due to finely divided gold.

Besides the borax bead and the charcoal stick, reduction is often
effected by means of metallic sodium, potassium, or magnesium.

Thus small amounts of phosphorus in anhydrous salts may be
detected in the following manner: The substance to be tested is placed
in a glass tube, 3 mm. wide and 50 mm. long, which is closed at one end.
A small cylinder of potassium or sodium (freed from petroleum by
rubbing between filter-paper), or even a piece of magnesium wire, is
added to the tube, and the contents then heated until the glass itself
begins te soften. The reaction is so violent that the substance seems to
take fire. After cooling, the tube is broken in a porcelain mortar,
when by breathing over the mass, the smell of phosphoretted hydrogen
may be detected.

68 GENERAL PRINCIPLES A,

The halogens, sulfur, and nitrogen are tested for in a similar
way, as will be shown later.

(d) Reduction in the Upper Reducing Flame for the Purpose of Forming
Metallic and Oxide Incrustations

The volatile elements which are reducible by means of hydrogen
or charcoal may be detected in this part of the flame with the great-
est ease, as, for example, arsenic, anti-
mony, cadmium, bismuth, selenium, and
tellurium. The metallic incrustations are
obtained by holding in one hand a small
portion of the substance on a thin asbestos
thread (platinum will be attacked) in the —
upper reducing zone of a small gas flame,
where the oxide is reduced to volatile
metal, and burned in the upper oxidizing
flame to oxide. In the other hand, closely
over the substance to be tested, is held
a glazed porcelain evaporating-dish, filled
with water, as is indicated in Fig. 3 at B.
The metallic vapors are condensed by

Fic, 3. the cold dish, and deposited on it in

the form of a metallic mirror or film.

If, however, the dish is held above the upper ‘oxidizing flame

(at A), there is formed a thin, often invisible, oxide incrustation on
the bottom.

Should it be necessary to treat the metallic incrustation with a
large amount of solvent (as is necessary in the detection of selenium
and tellurium), the porcelain dish is replaced by a test-tube half filled
with cold water. A somewhat larger test-tube is used to hold the sol- —
vent, and the smaller test-tube, on which the incrustation was deposited,
is placed within the larger tube and the liquid warmed if necessary.

(e) Blowpipe Reduction on Charcoal

These tests are made in the so-called “ preliminary examination.”
For this purpose a small cavity is made with a penknife in a piece
of good charcoal (preferably of linden wood), in which a knife-bladeful of
the substance to be tested is placed, previously mixed with twice as much
anhydrous sodium carbonate. As charcoal is a porous substance, it will
readily absorb melted substances, such as salt of the alkalies. Other

REACTIONS IN THE DRY WAY 69

substances are changed, by means of the sodium carbonate used, into
carbonates, which are, for the most part, decomposed, on heating, into
oxides and carbon dioxide. The oxides of the noble metals are decom-
posed, without the aid of the charcoal, into oxygen and metal; while
those of the remaining metals are either reduced to metal or remain
unchanged. Thus CuO, PbO, BizOz, Sb203, SnOe, Fe203, NiO, and
CoO are reduced either to a fused metallic globule (Pb, Bi, Sb, Sn, Ag,
and Au), or to a sintered mass of metal (Cu), or to a glistening metallic
fragment (Fe, Ni, Co, Pt). The oxides of zinc, cadmium, and arsenic
do not give metallic globules, but are, however, easily reduced to metal.
These metals are so volatile that they are changed into vapors, and are
carried from the reducing zone of the flame into the oxidizing zone, where
they are changed into difficultly volatile oxides. These oxides, which
have characteristic colors, are then deposited on the charcoal outside
the cavity.

Zine gives an incrustation which is yellow while hot, and white
when cold; that of cadmium is brown; while the oxide of arsenic gives
a white and readily volatile incrustation. Furthermore, the volatiliza-
tion of arsenic gives rise to a characteristic garlic odor. The metals
lead, bismuth, and tin give, besides the metallic globule, an oxide
incrustation which is typical.

At the same time, nitrates, nitrites, chlorates, etc., may be recognized
by the fact that they cause a very rapid combustion of the glowing —
charcoal (deflagration). This deflagration is not to be confused with
a decrepitation which takes place on heating substances containing
enclosed moisture or gases, such as rock salt, fluor-spar, etc. Crystals
of such substances are burst by the quick eee of the enclosed
liquid, and scattered about.

Many difficultly fusible substances do not melt into the charcoal.
Thus many silicates form a bead with the soda, which only after con-
tinuous heating will give up the alkali and allow it to be absorbed by
the charcoal, leaving behind the white infusible silica. Phosphates
and borates act similarly, only these do not leave behind an oxide,
but a fused glass. Infusible white oxides, as those of calcium, stron-
tium, magnesium, aluminium, and many of’the rare earths (Welsbach
mantle, for example), glow very brightly, and in fact more brightly
as they are more strongly heated.

70 GENERAL PRINCIPLES

Division of the Metals into Groups

The metals, for purposes of analytical chemistry, may be divided
into five groups:

The First Group contains those metals whose chlorides are insoluble,
or difficultly soluble, and whose sulfides are insoluble in dilute acids.
‘They may, therefore, be precipitated from their solutions by means
of either hydrochloric acid or hydrogen sulfide.

The Second Group contains those metals whose chlorides are soluble,
but whose sulfides are insoluble in dilute acids. They may be precipi-
tated from their solutions by means of hydrogen sulfide, but not by
hydrochloric acid.

The Third Group contains those metals whose sulfides are solu-
ble in dilute acids, but are insoluble in water and alkalies; and-also
those metals whose sulfides are hydrolytically decoupanet into
hydrogen sulfide and insoluble hydroxide. The members of this group
are precipitated completely by hydrogen sulfide only from alkaline
solutions.

The Fourth Geos contains those metals whose sulfides are soluble
in water, but whose carbonates are insoluble in the presence of ammonium
chloride. They are precipitated by ammonium carbonate in the presence
of ammonium chloride, but not by any of the above reagents.

The Fifth Growp contains magnesium and the alkalies; they are
not precipitated by any of the above reagents.

In order to carry out an analysis with certainty it is necessary
to understand not only the reactions of the different elements, but
we must know as well the sensitiveness of each reaction. The analyst
should be able to draw a conclusion by the size of the precipitate formed
as to the approximate amount which is present in the original substance.
This, however, is possible only when the experiments are made with
known amounts. Consequently reagents of a known strength are used
and allowed to act on known amounts of the different substances. Ac-
cording to the suggestion of R. Blochmann* it is well to make the solu-:
tions of the different reagents either double-normal, normal, half-
normal, or tenth-normal. For many years the author has used in his
laboratory solutions of reagents and salts according to this principle,

* Berichte, 1890, 31.

THE LABORATORY REAGENTS 71

and has found that the beginner in this way gets a far better under-
standing of the stoichiometrical relations than when solutions of almost
any concentration are used, as was formerly the custom.

By a normal solution is understood one which contains in a
liter one gram-equivalent of the substance in question, referred to a
gram atom of hydrogen as a unit. A tenth-normal solution will contain
one-tenth of a gram-equivalent in a liter, ete.

- Thus one liter of a normal solution will contain

HCl = 36.46 gms.
H2SO4 im 98.08 _ 49 o4 ems.
2 2
Ere 08-08 55’ 6g gms.
3 3 equivalent to one
NaOH = 40.06 gms. gram-atom of hydrogen.
KMn0s _ 158.11 _ 5) 69 gms.
5 5
Bor _ AE = 49.08 gms.

The great advantage of this system is that one always knows how
much of one solution should be used in order to react with another
quantitatively. Thus 1 cc. of a normal caustic soda solution will
neutralize 1 cc. of any normal acid, or 2 cc. of any half-normal acid.
In the same way 1 cc. of a normal solution of sulfuric acid, or of any
sulfate, will precipitate quantitatively the barium from 1 cc. of a normal
barium chloride solution.

The Laboratory Reagents

I. CoNcENTRATED ACIDS

Sp.gr. Per cent by Wt.| Approx. Conc.
MRMINOCTUOEIC BON oe dcocp.b 06 5 aces bess 1.19 37.9 12 N.
MMA LIT SO ONO 06 109 5 a PE AO ue dis een ern eee’ vat, Tt cate etnintets 24.N.
MN IMR N Shas ashe bo slnk ee FR acy ales o's 1.42 69.8 16 N.
MMIIOTIO BOWE oc.) bse ke we ss eles ee © 0 hoe 85 15.N.
Sulfuric acid........ eats ey OLY 1S Re gee 1.84 96.0 36 N.

* This is considering the solution as an oxidizing agent, As a precipitant,
the normal solution of KMnQ, would contain one mole per liter. The reagent
is almost inyariably used as an oxidizing agent and not as a precipitant.

72 GENERAL PRINCIPLES
II. Diturep Acips
Sp.gr. Per cent by Wt.} Approx. Conc. —
POCO BOG... |: . sce has Scalise 1.04 34.6 6N.
Hydrochloric acid 5.00 80... ee eee 1.10 20.0 6N.
PRISE OIE oe. Ca Es aka «bd Galas alee 1.20 32.3 6 N.
PePGHOriG Hid © 0. oe ee PS 1.12 tinea ‘2N.
SEO MOME Ooi 5+. dies beets ato lee aes 1.18 24.8 6N.
PEUPIIUONIS MONG oy. fs Soy, oie eis nd (Satd. soln. of SO2) 0.33 N.
SRPERPIO MIDI Se wicca oe wee (150 g. per liter) 2N.°
III. Basss
Sp.gr Per cent by Wt. Approx. Conc.
Ammonium hydroxide................ 0.90 28% NHs 15 N Ys
Ammonium hydroxide...........:.... 0.96 9.9%NHs 6 N.
Barium hydroxide. .................:. (satd. soln. Ba(OH)2-8H2O) 0.4 N.
Potassium hydroxide.................. 1.37 36.9 9 N.
EIN DVOTORIGC as ses cae 1,22 19.7 6 N.
IV. Sauts ©
(a) Ammonium Salts
Name of Salt. Directions for Making Reagent. Cone
Ammonium acetate Mix equal volumes of 6-normal acetic acid 3.N.
NH,C2H;02 or NH,Ac and 6-normal ammonia.
Ammonium carbonate | Dissolve250gms.freshly powderedammonium 6N.
(NH,)2COs carbonate in 1 liter 6-normal NH,OH. The
commercial salt is NH4sHCO;3+NH,CO.N Ha.
Ammonium chloride | Dissolve 54 g. in 1 liter of water. N.
NH.Cl
Ammonium molybdate | Dissolve 75 g. of pure ammonium molybdate} 0.85 N.
(NH,)2MoO, in 500 cc. water, pour the solution into 500
cc. of 6-normal- HNO; and shake till the
precipitate dissolves. “The formula of the
commercial salt is (NH4)sMo7O.4-4H20.
Ammonium oxalate | Dissolve 35 g. (NH4)2C20,-H2O in 1000 cc. | 0.5 N.
(NH4)2C204 water.
Ammonium polysulfide | Digest 1 liter 6-normal ammonium monosul- 6N.
(NH,4)2Sz fide with 25 gms. flowers of sulfur for some
hours and filter.
Ammonium sulfide Pass H.S into 200 ec. 15-normal NH,OH in 6N.
(NH,).S a bottle immersed in running water or in
ice water until no more gas is absorbed;
then add 200 cc. 15-normal NH,OH and
dilute to 1 liter. ;

THE LABORATORY REAGENTS 73
(b) Other Salts
Name of Salt. Formula. Wooue Nest ead Cone

EE RSE Nhs role aa» ao KAI(SO4)2-12H20...... 475 59 0.5 N.
Barium chloride........ BaCh:-2H.0.....5....0. 244 122 N.
Bismuth nitrate........ Bi(NOs3)3-5H2O0........ 484 80 0.5 N.
Cadmium nitrate....... Cd(NO3)2°4H2O....... 308 154 N.
Cadmium sulfate....... 38CdS0O,4-8H20......... 770 64 0.5 .N.
Calcium chloride....... CaCl, -6H20........... 219 110 N.
Calcium sulfate........ AMM IAM LEM. bere s 2 oboe 172 2.6 Satd.
Chrome-alum.......... KCr(SO,4)2-12H2O...... 399 83 0.5 N
Cobalt nitrate........:. Co(NOs3)2:6H20........ 291 50 1% Co
Copper sulfate......... CuSO,-5H20.......2.. 250 125 N.
Ferric chloride......... POC GH i odes ss 270 90 i a
Lead acetate........... Pb(C2H302)2-3H20..... 379 190 N.
Magnesium sulfate. .... MgSO,-7H20.......... 247 123 N.
Manganese sulfate...... MnS0O,-4H20.......... 223 56 0.5 N.
Mercuric chloride....... POCA as gaia se Sand 8 ow ah 272 27 0.2 N.
Mercurous nitrate.... ~. Hg2(NOs)e.......--.-. §25 263 N.
Nickel sulfate.......... NiSO,-7H20.......... 281 70 0.5 N.

, Potassium bichromate.. .| K;Cr.O7.............. 294 49 N.f

- Potassium chromate... .| K,CrO,............... 194 291 3.N.?
Potassium cyanide...... FECA ee ats on Din aus ys 65 65 N.
Potassium ferricyanide. .| K3Fe(CN)............. 329 110 N.*
Potassium ferrocyanide..| K,Fe(CN).s-3H2O...... 422 105 N.*
Potassium iodide....... 1 10 Ma ee ere es geen ae mk heeaien 166 17 0.1 N.
Potassium nitrite... .... BEV ee St kare ee elak 85 255 3 N.
Potassium permanganate} KMnO,.:............. 158 10 0.3 N.f
Potassium thiocyanate. .| KCNS................ 97 97 aS.
Silver nitrate.......... WISIN erectus peleacate b 170 17 0.1 N.
Sodium acetate......... NaC2H;02-3H20....... 136 136 N.
Sodium arsenite........ PRG ics oes ehe ek 130 130 N.
Sodium bromide........ NaBr-2H,0 or NaBr.. ./139 or 103} 69 or 51 | O.5 N.
Sodium carbonate...... INN ac heie nals Dees 7 106 159 3 N.
Sodium hypochlorite....| NaClO............... 74 37 N.f
Sodium nitrite......... INGINONS Sais Fate Weenie} «8 2 69 207 3N.
Sodium phosphate...... NaeHPO,:12H20....... 358 119 N.
Sodium thiosulfate...... NaeS.0;-5H20......... 248 124 0.5 N.f
Uranyl acetate. .4..../1 UO2(C2H302)2+2H20. ake 425 21 0.1 N.

* As precipitant.

Pt) in 100 ce. water.

t As oxidizing agent.

ft As reducing agent.

V. SpeciAL REAGENTS
Chloroplatinic acid: Dissolve 26.5 gms. H2PtCl,-6H2O (corresponding to 10 gms.

Kther saturated with HCl: Saturate anhydrous ether at 0° with dry HCl gas.
Ferrous sulfate, 2 N: Dissolve 280 gms. FeSO,-7H2O in 6 N oa acid _and keep
in contact with iron nails.

Hydrogen peroxide, 3 per cent.

fre

User RA

Magnesium ammonium chloride, N in MgCl: ‘Dissolve 100 gms. MgCl, -6H2O and
100 gms. NH,Cl in water, er, add 50 cc. 15-normal NH,OH and dilute to one liter...

Potassium mercuric iodide, 0.5 normal in K,HglI,: Dissolve 115 gms. HglIy and 80
gms. KI in enough water to make the volume 500 cc.; add 500 ce. 6-normal
NaOH and decant the solution from any precipitate that may form on standing.

used, as well as the solids, must be absolutely free from ammonia.

The stock solution should be kept in the dark. For sensitive work the water

74 GENERAL PRINCIPLES

Potassium pyroantimonate: Add 20 gms. of the best commercial salt to 1 liter of
boiling water, boil until nearly all the salt has dissolved, cool quickly, add 30
ec. of 10 per cent KOH solution, and filter. :

Sodium cobaltinitrite: Dissolve 250 gms. NaNO: in 500 ce. water, add 150 ce.
6-normal HC,H;O2 and 25 gms. Co(NO;3).-6H2O. Let the mixture stand over
night, filter and dilute to 1 liter.

Stannous chloride, l-normal: Dissolve 113 gms. SnCl.-2H:O in 100 cc. 12-normal
HCl, dilute to 1 liter and keep in bottles containing a strip of pure tin.

Starch and potassium iodide: Rub 20 gms. soluble starch to a thin paste with.
a little water in a mortar and pour the paste into 1 liter of boiling water.
Boil five minutes and filter through a loose plug of cotton wool. Add 10 gms.
KI and 5 ce. chloroform to the filtrate. ~

Turmeric: Shake turmeric powder with 95 per cent alcohol and filter.

Urea: Dissolve 200 gms. urea in 1 liter 6-normal HCl.

VII. SatTuraTeD SOLUTIONS

Barium hydroxide ................
Bromine water). F6 Ase
SO oe UE eh eee a
Chiorme water: . 60044 S.icc
Hydrogen sulfide..................
BAVC WRU i sk bos coo we Gtr ees

'a\e + oe © bie #

1000 cc. water dissolves.
68 gms. Ba(OH)2 at 20°.
32.68 gms. Bre.

2.6 gms. CaSO,-2H,0.
6.5 gms. Cle.

VII. Sprcrau SoLvENtTS

Amy] alcohol.
Carbon disulfide.
Chloroform.

Ethyl alcohol (95 per cent).
Methyl alcohol (free from acetone).

VII. Souip REAGENTS

Absorbent cotton

Ammonium chloride, NH,Cl
Ammonium nitrate, NH,NOs3
Barium carbonate, BaCOs

Borax, NazB,O;

Boric acid, HBO;

Calcium chloride (anhydrous)
Calcium oxide

Calcium phosphate

Chloride of lime, CaOCl-Cl
Copper wire or turnings

Ferrous sulfate, FeSO, -7H,O
Ferrous sulfide, FeS :
Iron nails, Fe

Lead (finely granulated) Pb

Lead dioxide (free from Mn), PbO:
Litmus paper, blue

Litmus paper, red

Paraffin

Potassium acid sulfate (fused), KHSO,
Potassium carbonate, K,COs

Potassium chlorate, KCIO;

Potassium dichromate, K2Cr,O,7

Potassium ferricyanide, K;Fe(CN)g

Potassium iodide, KI

Potassium nitrate, KNO;

Silica (precipitated), SiOz

Silver sulfate, AgeSO,

Sodium acetate, NaC.H;0,

Sodium ammonium phosphate (micro
cosmic salt), NaNH,HPO,-4H,O

Sodium bismuthate, NaBiO;

Sodium carbonate, NazCOs

Sodium peroxide, Na,Oz2

Sodium nitroprussiate,
Nazk’e(CN)s ° NO -2H.O

Sodium sulfide, NaS -9H:O

Sodium tetraborate, (see Borax)

Starch, (CeHiOs)~

Tartaric acid, H2CsH,O¢

Tin (finely granulated)

Zine (finely granulated)

DETERMINATION OF THE SENSITIVENESS OF REACTIONS 75

Determination of the Sensitiveness of Reactions.

| The more sensitive a reaction is, the smaller will be the amount
_ of the substance which can be detected in a given volume, in a definite
time, with the reagent in question. Let us assume that the amount
of substance taken is dissolved in 100 cc. of liquid, the time allowed:
- to be two or three minutes, and the limit of sensitiveness to be the
smallest amount of substance which can be detected under these con-
ditions.

A few examples will make the method clear:

Magnesium salts are precipitated by means of sodium phosphate,
in the presence of ammonium chloride and ammonia, in the form of
magnesium ammonium phosphate. What is the sensitiveness of this
reaction? We take 1 cc. of our normal magnesium sulfate solution,
add three drops of ammonium chloride solution, and two or three
drops of ammonia and sodium phosphate solutions; the characteristic
white precipitate is formed immediately. We dilute, now, the normal
solution of magnesium sulphate ten times, and repeat the experiment
with 1 cc. of the diluted solution. The result will be—

1 ce. of N. Mg solution, 100 cc. =1.2 g. Mg, reacts immediately.

mee. .-N/10;: Mg “ 100 ec. =0.12 “Mg,

icc. N/100 Mg e 100. cc: =0:012. -- “* Mg, ~. * s

lec... N/1000 Mg 100 cc.=0.0012 “ Mg, ‘* afterafewseconds.

lec. N/10000 Mg - 100 cc. =0.00012 ‘ Mg, “ after one or two
minutes.

If, therefore, 100 cc. of a solution contain 0.00012 gm. Mg, the
magnesium can be detected within one or two minutes. Should the
detection of smaller amounts be desired, the solution must be con-
centrated by evaporation.

This reaction can be called very sensitive. The following potassium
reactions are much less delicate:

(a) Reaction with Chloroplatinic Acid (page 78)

1 cc. of 0.2 N KCl solution, 100 cc. =0.78 gm. K, reacts with a drop
of H2PtCle immediately.

1 cc. of 0.02 N KCl solution, 100 cc.=0.078 gm. K, does not cause
precipitation within three minutes. .

1 cc. of 0.04 N KCl solution, 100 cc. =0.156 gm. K, does ee cause
precipitation within three minutes; but does, however, on addition
of two drops of alcohol.

1 ce. of 0.06 N KCl solution, 100 cc.=0.234 gm. K, reacts imme-
‘diately on stirring.

76 GENERAL PRINCIPLES

The sensitiveness of the reaction lies, therefore, between 0.156 and
0.234 gm. K per 100 cc. In order to detect smaller amounts of potas-
sium than 0.156-0.234 gm. per 100 cc., the solution must be strongly
concentrated by evaporation

(b) Reaction with Tartaric Acid (page 79)

1 cc. of 0.2 N KCl solution, 100 cc.=0.78 gm. K, reacts immedi-
ately with two drops of sodium acetate and two drops of a concentrated °
solution of tartaric acid.

1 cc. of 0.02 N KCI solution, 100 cc.=0.078 gm. K, reacts after

one to two minutes with vigorous shaking.

This can be taken as the limit of sensitiveness.

If the beginner will test the delicacy of reactions in this way, he
will quickly get a clear insight into the solubility relations of the
different salts.

Probably the quickest way of learning these relations is to analyze
first of all solutions known to contain an equal quantity of each constit-
uent of a given group. It is well for the beginner to start with 100.
ec. of solution and 0.1 gram of each dissolved cation or anion.

When the sensitiveness of the reactions is known, qualitative
tests are often the most accurate methods for estimating small quantities
. of substances. Thus it is possible to detect very small quantities
of titanium by means of the hydrogen peroxide test, and the most —
accurate method for determining such small quantities is to compare
the intensity of the color change with that produced in a series of —
solutions containing known quantities of titanium.

PART II. REACTIONS OF THE METALS (CATIONS)

The separation of the metals into groups (cf. p. 70) is based upon
the varying solubilities of the chlorides, sulfides, hydroxides and car-
bonates. In general, the metals removed first in the scheme of analysis
are those which form the smallest number of soluble salts and the metals
tested for last are those which form the largest number of soluble salts.
In qualitative analysis, the relatively insoluble salts occupy the chief
interest and it is, therefore, simplest to begin the study of the subject

with that group of metals which is tested for last. Compounds con-

taining these metals are very common and are used to some extent in
the separation and identification of the metals of the other groups. A
knowledge of the characteristic reactions of such compounds is neces-
sary in order to understand the chemistry involved in the analysis of
the other groups, and this furnishes another important reason for
taking up the study in the reverse order to that in which the analysis
is usually carried out. In this part of thé book only the reactions of
the more common elements will be considered. The rarer elements,
such as lithium, cesium and rubidium of the alkali group, will be con-
sidered in Part V.

GROUP V. THE ALKALI GROUP

POTASSIUM, K; SODIUM, Na; and AMMONIUM, NH4

The metals potassium and sodium are the most reactive of all the
common positive elements (cf. p. 40). They oxidize very rapidly when
exposed to the air and decompose water at ordinary temperatures; the
hydrogen of water is reduced to the gaseous condition and the metal
is oxidized to alkali hydroxide in equilibrium with alkali cations and
hydroxyl anions. On account of the extent of the ionization (cf. p. 10)
the alkali hydroxides form very strong bases. The solid hydroxides
are the most stable of all hydroxides; they do not break down into oxide

and water even on being melted. The pure oxides are difficult to pre-

pare; cautious heating of the metals in air results in the formation of '
considerable peroxide as well as oxide.
77

78 REACTIONS OF THE METALS

Ammonium is classed with the alkali metals because the solubility ©
of ammonium salts is similar to that of potassium salts. The ammonium
radical differs from the alkali metals in being capable of oxidation, the
usual product of the oxidation being nitrogen gas. Ammonium hydrox-
ide is a much weaker base than the other hydroxides, and the salts are
either volatile or are decomposed on being heated. The salts of potas-
sium, sodium and ammonium are colorless for the most part, and
readily soluble in water. Of these salts the carbonates, the tertiary
and secondary phosphates, the cyanides, and the borates react alkaline in —
aqueous solution (hydrolysis). The salts of the alkalies are more or
less volatile and impart to the non-luminous flame characteristic colors.
When a solid is involved, either as initial substance or as final product,
in any of the characteristic reactions of this and the following groups,
it will be designated usually by bold-faced type in the equation. The
formation of a gas will usually be indicated by placing an arrow after
the symbol in an equation.

POTASSIUM, K. At. Wt. 39.10
Sp. Gr. 0.87. M. Pt. (Melting-point) 62.5° C.

Occurrence.—Sylvite (KCl), isometric, and carnallite (MgCl,
KCl-6H2O) orthorhombic, occur at Stassfurt in the presence of halite
and anhydrite. Saltpetre (KNOs), orthorhombic prisms. Further,
in very many silicates, e.g., monoclinic feldspar (KAISisOg), and
muscovite (KHe2AlsSigOi2); also in plants in the form of organic
salts, which yield on combustion potassium carbonate (potash).

REACTIONS IN THE WET WAY

| Potassium forms very few salts that are difficultly soluble in water.
The chloroplatinate, acid tartrate, and perchlorate are the least soluble,
and are consequently used in the detection of potassium.
1. Chloroplatinic Acid,* Hy2{PtCl¢|], gives in concentrated solutions
of the chloride a yellow precipitate of potassium chloroplatinate,

PtCh-"+2Kt—KoPtCly,

which consists of small regular octahedra (visible with a magnify-
ing-glass). In case. the potassium solution is not very concentrated,
no precipitation may appear at first; but on rubbing the sides of the

* Platinic chloride, PtCl, gives no precipitate with potassium salts, or at least
only after long standing. The above reagent, chloroplatinic acid, is a dibasie acid
and is obtained by dissolving platinum in aqua regia. The solution is prepared of
such strength that there are 10 gms. of platinum in every 100 ¢e.

POTASSIUM | 79

beaker or test-tube with a glass rod the formation of the precipitate
will be hastened.

This is always the case when a crystalline precipitate is formed.
The solution is supersaturated before the precipitate separates out,
_ and the formation of crystals is hastened by the mechanical shock.

The behavior of the potassium chloroplatinate on ignition is charac-
teristic; it is decomposed into chlorine, platinum, and potassium chloride:

K2[PtClo] =2KC1+Pt+2Cle 1.

If the products of ignition are treated with water, and the platinum
filtered off, the filtrate will again give with chloroplatinic acid the yellow
erystalline precipitate of K2[PtCle]. (Note difference from ammonium
chloroplatinate.)

Solubility of the Potassium Chloroplatinate in Water.

100 cc. of water dissolve at 0°, 0.70 gm.; 10°, 0.90 gm.; 20°, 1.12
gms.; and at 100°, 5.18 gms. Ke[PtCle].

In a saturated KCl solution, or in 75 per cent alcohol, the pre-
cipitate is practically insoluble.

For this reaction it is best to use the chloride. The addition of
chloroplatinic acid to potassium iodide solution causes a deep-reddish-
brown color due to the conversion of PtCls~ into non-ionized PtI¢~ -
ion, of which the potassium salt is more soluble:

PtCle~ +6I “—> PtI,— +6Cl~.

Similarly, potassium cyanide is not precipitated by chloroplatinic
acid, owing to the formation of complex platinum-cyanogen compounds.

To test an iodide or cyanide for potassium, the salt should first be
changed to chloride by evaporation with concentrated hydrochloric acid.

2. Tartaric Acid, H2C4H40¢, produces, in not too dilute neutral
solutions of potassium salts, a white crystalline precipitate of potas-
sium acid tartrate (orthorhombic, hemihedral) :

K*+H2C4Hs40g—-KHCsHi0¢ +H,

Rubbing the sides of the dish will hasten the formation of the
precipitate. |

Potassium acid tartrate is readily soluble in mineral acids, but
difficultly soluble in acetic acid and water; 100 parts of water at
10° C. dissolve 0.425 gm. of this salt. If sodium acetate is added to
the solution, the hydrogen ions set free by the above reaction will
unite with the acetate ions to form non-ionized acetic acid (cf. p. 46):

H*+C2H3027 > HC2H302,

80 REACTIONS OF THE METALS

whereby the reaction is made much more sensitive. 'Too much sodium
acetate, however, exerts a solvent action upon the potassium acid tar-
trate. In the presence of considerable sodium acetate, acetic acid is.
ionized to such a slight extent that the hydrogen ions from HC4,H40,—
disappear and, as a result, the potassium acid tartrate dissolves. Neu-
tralization with caustic alkali is even more dangerous; soluble alee
tartrate is formed:

KHC,H.,0,+0OH — Kr+ C4H40, +H20.

On igniting potassium acid tartrate, empyreumatic vapors (smelling
like burnt sugar) are given off and a mixture of carbon and potassium
carbonate is left behind. If the mass is now moistened with hydrochloric
acid, it will froth strongly. This is a property not only of potassium
tartrate, but is common to all alkali salts of organic acids. On ignition
they are changed into carbonates, and when the acid is non-volatile,
carbonization takes place; but with volatile acids there is at the most
only a slight carbonization. With many metals the carbonate is not
left unchanged; frequently it is broken up into carbon dioxide and
oxide of the metal; in the case of salts of organic acids with reducible
metals, the metal itself is left with the carbon. Thus sodium acetate
will yield sodium carbonate and acetone, with only a slight carboniza-
tion:
2CH3CO -ONa= NaeCO3 +CHsCOCHs.

Acetone (colorless
volatile liquid,
burns with
luminous flame)
On gentle ignition, calcium oxalate yields calcium carbonate and

carbon monoxide; the latter burns with a blue flame.

CaC204=CaCO3+COT.

On strong ignition, the calcium carbonate is decomposed into lime and
carbon dioxide:

CaCO; — CaO-+-COs; fT.

Tartrates of lead, iron, and many other metals on being ignited yield
carbon and metal.

3. Bismuth-Sodium Thiosulphate (Carnot’s* reaction).—If one drop
of half-normal bismuth nitrate solution is mixed with two or three
drops of half-normal thiosulfate solution and 10-15 cc. of absolute
alcohol (any turbidity being removed by the careful addition of a very

*Z. anal. Chem. (1897), 512.

ras.

POTASSIUM 81

little water), a sensitive reagent for potassium ions is prepared in which
the bismuth is present in a complex anion:

Bit **+38203" — Bi(S20s)3".

The sodium salt is soluble in alcohol but the yellow potassium salt is not:

Bi(S203)3-+3K* — K3[Bi(S203)3).

The presence of ammonium chloride prevents the reaction.
4, Fluosilicic Acid, H2Sil’s, added in considerable excess to a solution
of a potassium salt, precipitates gelatinous potassium fluosilicate,

Sik’. + QKt = K2SiF¢ ’
which is difficultly soluble in water and dilute acids and insoluble in

alcohol. On heating, it is decomposed into volatile silicon fluoride,
and potassium fluoride remains behind:

K.2SiFs=2KF+SiF; 7.
5. Perchloric Acid, HC1O4, precipitates white, crystalline potassium

perchlorate,

HClO. +K+t— Ht+KCl0,;

100 ce. of water at 0° dissolve 0.07 gm., and at 100° 19.8 gms. KCI1Ou.
It is so slightly soluble in 97 per cent alcohol that the precipitate can
be obtained with less than 2 mgs. of potassium ions.

6. Sodium Cobaltinitrite, Nas3[Co(NOz)«], precipitates yellow
potassium-sodium cobaltinitrite from neutral or slightly acid solution:

Nas[Co(NO2)¢] +2KCI=K2Na[Co(NO2e).¢]+2NaCl.

Ammonium salts give a similar precipitate, but moderate amounts of

~alkaline-earth elements or of lithium and sodium do not interfere. -

The test must not be made in alkaline solution or Co(OH)3 will be pre-
cipitated. The reagent, prepared according to the directions on p. 74,
permits the detection of 0.3 mg. of potassium within ten minutes. If
the reagent is prepared according to the following directions of Biil-
mann,* as little as 0.0009 mg. potassium can be detected in the presence
of 4000 equivalents of sodium.

PREPARATION OF SODIUM COBALTINITRITE

Dissolve 150 gms. of sodium nitrite in 150 cc. of hot water and allow the
solution to cool to about 40°, which will cause the deposition of some sodium
nitrite crystals. Add 50 gms. of cobalt nitrate crystals, stir rapidly and add
50 ce. of 50 per cent acetic acid in small portions; stopper the flask and shake
vigorously. Pass a rapid stream of air through the liquid and then allow it

* Z. anal. Chem., 39 (1900), 284.

82 REACTIONS OF THE METALS

to stand quietly over night. In the morning, more or less brown precipitate
will be found on the bottom of the flask due to a little potassium salt present
in the sodium nitrite. Siphon off the clear liquid through a filter and add to
the filtrate, while stirring, about 200 cc. of aleohol in small portions; this causes
the precipitation of the greater part of the dissolved sodium cobaltinitrite.
After several hours, filter off the precipitate and drain it as completely as possi-
ble with the aid of suction. Wash the precipitate four times with 25-ce. portions
of alcohol and twice with ether. Recrystallize the salt by dissolving each 10
gms. of solid in 15 ec. of water and precipitating with 35 cc. of alcohol. The dry
salt keeps very well, but the aqueous solution decomposes gradually. To
obtain the best results, the reagent should be freshly prepared by dissolving
1 gm. of the salt in 10 ce, of water. .

REACTIONS IN THE DRY WAY

Potassium compounds color the non-luminous flame’ violet. The
presence of very small amounts of sodium obscures the violet color,
but if the flame is viewed through cobalt glass or indigo solution, the
reddish-violet potassium rays pass through, while the yellow sodium
rays are completely absorbed. |

Flame Spectrum.—Potassium gives a characteristic flame spec-
trum. A double red line, 769.9 yu and 766.5 yy (appearing as a single
line with weaker dispersion), and a faint violet line, 404.4 uu, appear
at comparatively low temperatures. With a hotter flame, other
lines are visible in the yellow, 583.2 yu; 580.2 wu and 578.2 wy, and in »
the green, 535.1 wu and 511.3 py (see chart, Frontispiece).

Gooch and Hart * were able to detect 0.001 mg. of pure potassium
as chloride in 0.02 cc. water, but the presence of one hundred times
as much sodium caused the potassium lines to disappear unless the
sodium rays were deflected from the field.

SODIUM, Na. At. Wt. 23.00
Sp. Gr. 0.97. M. Pt. 95.6° C.

Occurrence.—Sodium occurs very extensively in nature. Its most
important mineral is halite, rock salt (NaCl), isometric system.
Halite is found in very large deposits often quite pure, but usually
contaminated with clay, anhydrite, and gypsum, and is present —
in large amounts in the ocean, and in many salt springs. Sodium
also occurs in nature in the form of carbonate, as thermonatrite
(Na2CO3-H20), orthorhombic; natron or soda (Na2CO3-10H20),
monoclinic; trona (NazCOz3-NaHCO3-2H2O), monoclinic; as nitrate |
in Chili saltpetre, or soda nitre (NaNOs3), hexagonal, rhombohedral;
as eryolite (NagAlF’s), triclinic; in many silicates as albite (NaAISi3Qg),
triclinic; and as tinkal, borax (NazBsO7-10H20), monoclinic.

* Z. anal. Chem., 36 (1897), 390.

SODIUM 83

REACTIONS IN THE WET WAY

1. Potassium Pyroantimonate,* KsH2Sbe207, produces in neutral
or weakly alkaline solutions of sodium salts a heavy, white, crystalline
precipitate, which is formed more quickly by rubbing the sides of the ,
vessel with a glass rod:

KeHeSbe07 + 2NaCl =. a2HeSbeO07 + 2KCI.

The test must not be made in an acid solution, for in that case
an amorphous precipitate of pyroantimonic acid will be formed:

KeH2Sb207+2HCl = H4Sb207+2KCl.

Furthermore, no other metals than the alkalies sheuld be present,
because they also cause precipitates—amorphous ones for the most
part.

2. Tartaric Acid and Chloroplatinic Acid do not precipitate
sodium salts, the sodium ‘salts of these acids being soluble in alco-
hol as well as in water. (Note difference from potassium.) Sodium
chloroplatinate is orange in color.

3. Hydrochloric Acid and Alcohol precipitate sodium chloride.
Sodium chloride is prepared pure for chemical purposes by passing
hydrogen chloride gas into the saturated aqueous solution of the salt
and expelling the moisture and hydrochloric acid from the crystals
by heating them. If dry hydrogen chloride is passed into an alcoholic
solution of a sodium salt, less than 1 mg. of sodium will remain in
solution. |

Sodium Peroxide, NazO2

This substance, which is now used commercially on account of
its energetic oxidizing power, is obtained as a heavy, yellow powder,
by burning dry sodium in the air; it shows the following character-
istic reactions:

Behavior toward Water.—If a little water is added to some of this
substance in a test-tube, considerable heat is evolved and oxygen gas
is liberated (sufficient to ignite a glowing splinter ¢). Water decom-
poses the sodium peroxide, according to the equation

Naz02+2H20 =2Na0H+Ho20o.

* For the preparation of this reagent see page 74 and under Antimony.

+ This will sometimes cause an explosion. Commercial sodium peroxide often
contains metallic sodium, which with water forms hydrogen; thus both hydrogen
and oxygen are set free at the same time, and the glowing splinter may then cause
an explosion. (Private communication from E. Constam.)

84 REACTIONS OF THE METALS

But on account of the heat of the reaction a part of the hydrogen —
peroxide is decomposed into water and oxygen.

If the solution is kept cold, which can be done by throwing the
sodium peroxide in small portions into ice-water, it will dissolve with
scarcely any evolution of oxygen, to a clear, strongly alkaline liquid,
which gives, as before, all the reactions of hydrogen peroxide.

If some sodium peroxide is placed on a watch-glass under a bell-
jar and near an evaporating-dish containing water, the sodium per-
oxide in twelve hours will completely change over to a pure white
hydrate (NagO2+8H20), which will dissolve in water without decom-
position at the ordinary temperature. By standing in a desiccator
over sulfuric acid, the octohydrate is changed to NazO2+2H:20.

Reactions of Hydrogen Peroxide

(a) In Acid Solution

If the solution obtained by the action of water on sodium per-
oxide is used for these tests, it must be acidified with dilute sulfuric
acid, care being taken to keaists the solution cool.

1. Titanium Sulfate gives a distinct yellow color, caused by
the formation of pertitanic acid,

4
hy t4 HeOo +. 2HeO <> 4H HeTiOs

This is the most delicate test for hydrogen peroxide. The tita-
nium sulfate solution for this reaction may be prepared by fusing
one part of commercial titanium dioxide with 15-20 parts of potassium
pyrosulfate and dissolving the fusion, after cooling, in cold, dilute
sulfuric acid. It may also be prepared by heating titanium dioxide
with concentrated sulfuric acid, until a clear solution is obtained,
cooling and diluting carefully i water.

The addition of caustic alkali, ammonia or ammonium carbonate
gives a yellowish-orange precipitate which redissolves in an excess of the
reagent. Classen has used this reaction as a method for separating
titanium from ferric iron.* Some chemists prefer to write the formula
of pertitanic acid as TiO2-H2O2, which assumes a true peroxide
structure instead of hexivalent titanium. Possibly a condition of
equilibrium exists between the two structures:

ou ee O-.OH
ae > O= TS

* Ber., 21 (1888), 370.

“3
J

;

>
a
J
:
y.
A

”

SODIUM | 85

2. Chromic Acid.—If the acid solution of hydrogen peroxide is
shaken with a little ether (free from alcohol) and a trace of potassium
dichromate is added, after which the mixture is again shaken, the
upper layer of ethereal solution will be colored a beautiful blue, owing
to the formation of chromium peroxide (cf. p. 18).

This test is very sensitive and will detect as little as one-tenth
milligram of H2O2. In carrying out this test, a blank test must always
be tried with the ether and alcohol alone, because the former will often
give the test. Ether, after standing in the air, is likely to contain
some ethyl peroxide (C2H5)403, (?) which gives the test. It is possible
to free the ether from this peroxide by letting it stand over night
in contact with sodium and then redistilling it.

3. Permanganic Acid in acid solution will be decolorized, with
evolution of oxygen:

2Mn04-+5H202+6Ht > 2Mn+++8H20+50> 1.

Similar to the permanganate, many other oxides are reduced by

hydrogen peroxide, with evolution of oxygen; e.g., AgeO, Pb3Qa,
PbO2, MnOzg, Co20s, ete.:

Ago0+H202— H20+02 T +2Ag
Mn02+H202+2H* > Mn**+2H20+02 T
Co203+H202+4H* — 2Cott+3H20+02 1.

4. Potassium Ferricyanide and Ferric Chloride.—If a trace of
potassium ferricyanide is added to a very dilute and nearly neutral
solution of ferric chloride, so that the solution appears a distinct
yellow, and a nearly neutral solution of hydrogen peroxide is then
added, the mixture will soon assume a green tint, and finally, on stand-
ing, Prussian blue will separate out. The potassium ferricyanide
is reduced by the hydrogen peroxide to potassium ferrocyanide,
which forms Prussian blue with the ferric chloride.

2Fe(CN)6~ +H202 — 2Fe(CN)g--+02 T +2H*
and 3Fe(CN)g _+4Fet** — Fea[Fe(CN)¢]3.

According to Schénbein, as little as 0.02 mg. H2Oz per liter may be
detected by this reaction. As many other substances (SnCle, SOs,

etc.) will reduce potassium ferricyanide to potassium ferrocyanide,

this reaction alone is not always a reliable test.

5. Starch Paste and Potassium Iodide.—If to an acid solution
containing starch paste and potassium iodide some hydrogen peroxide
is added, a blue color will at once appear:

2K1+H202=2KOH-+Iz.

86 REACTIONS OF THE METALS

By means of this reaction, 0.05 mg. per liter of hydrogen peroxide may
be detected.

(b) In Alkaline Solution

1. Gold Chloride by means of hydrogen peroxide at ordinary
temperatures will be reduced to metal, with evolution of oxygen.
The gold usually separates in a very finely divided state, and appears
brown by reflected light and greenish blue by transmitted light:

2Autt+++3H.0.+60H~ > 2Au+6H20+302 fT.

If very dilute gold solutions are used, the metal sometimes separates
out in the form of a yellowish film adhering to the sides of the test-tube.
2. Salts of Manganese and Cobalt give dark-colored precipitates:
Mntt+20H~+H:202 — H20+Mn0O(0H)>2;

Brown

2Cot *+20H +H202 — H20+2Co(OH)3.
Black
Hypochlorites give the same reactions with manganese and cobalt
salts, but they do not give the reaction with gold chloride.

Ozone, O3

Ozone is always formed when oxygen is exposed to the silent electric
discharge. It is often present in oxygen that has been prepared elec-
trolytically and, according to Brunck, is present to:some extent in the
gas prepared by ignition of potassium chlorate. Ozone is a strong
oxidizing agent and behaves in many respects like hydrogen peroxide,
with which it is often confused. Ozone may be distinguished from
hydrogen peroxide as follows:

1. Ozone does not give a yellow coloration with titanium sulfate
solution. .

2. Ozone does not cause precipitation of gold from its solutions.

3. Ozone sets free iodine immediately from dilute, neutral potassium
iodide solution.

4. Ozone liberates bromine from an acid solution of sodium bromide.

5. Ozone causes bright metallic silver at once to assume a steel-
blue tint.

The sensitiveness of this last reaction is remarkable if carried out
according to the directions of Manchot and Kampeschulte. Heat a
bright piece of silver foil to about 240° and then expose it to the action
of ozone; steel-blue spots with violet edges at once appear. This
reaction does not take place with pure silver in the cold. If, however,
the silver is polished by rubbing with emery paper, the reaction will

r Pinta al >
vy at

AMMONIUM | 87

, then take place in the cold. Traces of iron oxide are left upon the
silver from the emery and catalyze the reaction. Other oxides, AgoO,

CoO, NiO, BizO3, PbsO04, -V205s, MnOz, CuO, ThO2, CeOe, TiOse,
WO3, U30s, and to a less degree MoOz, HgO, CaO and BaO, have a
similar effect. Thus if the silver is etched with nitric acid and is then

_ dried, it will react with ozone in the cold.

The principal reactions of sodium are the

REACTIONS IN THE DRY WAY

Sodium salts color the non-luminous gas-flame a monochromatic

_ yellow, which can be readily distinguished from the yellow flame of
the gas in the following way: If we illuminate an orange-colored body

(such as a stick of sealing-wax or a crystal of potassium dichromate)
with white light (all glowing solid bodies emit white light), the red and
orange rays will be reflected: the body appears orange. If these bodies
are illuminated with the monochromatic sodium light, they can now
only reflect yellow light: the bodies appear yellow (a delicate test).

Flame Spectrum.—A yellow double line (589.6 wu and 587.0),
coinciding with the D-line of the sun’s spectrum. This is an extremely
delicate reaction; 110-7 mg. of sodium can be recognized in the spec-
trum.

AMMONIUM, NHy. At. Wt. 18.04

Occurrence.—In small amounts as carbonate and nitrite in the
air; as ammonium chloride in the fissures of active volcanoes. Am-

‘monium derivatives are formed by the decay of many organic substances

containing nitrogen: albumin, urea, etc.,
CO(NHz)2+H20 =C02+2NHs,

and in a similar way by the dry distillation of many nitrogenous sub-
stances, such as coal, horn, hair, etc.

Although ammonium itself is known only in the form of its amalgam,
we are justified in considering it as a metal, in the first place because the
electrolysis of ammonium salts causes the setting free of the cation
NH,4(NH3+H) at the same time. that the corresponding anion is set
free; and, further, because the ammonium salts are isomorphous with
potassium salts.

REACTIONS IN THE WET WAY

ra: Strong bases, NaOH, KOH or Ca(OH)s, added to an ammonium
salt in the presence of a little water cause the evolution of ammonia
on heating; the gas can be recognized by its odor, by fumes of
ammonium chloride being formed when a rod moistened with 12 N

88 REACTIONS OF THE METALS

hydrochloric acid is placed in contact with the vapors, by its turning
red litmus blue; or by the blackening of mercurous nitrate paper:

Hg
2Hg2(NOs)2-+4NHs-+H20=3NHiNOs-+0C NH2-NO3+2Hg
Hg

Black

The reaction of strong bases upon ammonium salts may be explained on
the basis of the laws of chemical equilibrium. Ammonia, NH,, is a gas which
is very soluble in cold water and insoluble in boiling water. One volume of
water at 0° and 760 mm. dissolves 1800 volumes of the gas; at 20° it dissolves
710 volumes; at 100° all the gas can be expelled, there being no constant boil-
ing mixture as in the case of hydrochloric acid (p. 57).

The solution of ammonia in water at the laboratory temperature is in
equilibrium with ammonium hydroxide,

NH;+H:0 — NH,OH,

and, for this equilibrium, the mass action expression (p. 14) is
(N H;|X [HO]
[NH.OH]

In a dilute solution of ammonia, the absolute quantity of water present is not
changed much as a result of this equilibrium, so that when an accuracy corre-
sponding to only two significant figures is desired, the H,O member of this
expression may be regarded as a constant (cf. p. 48) and the equilibrium
expression then becomes

=k.

[NH3] _ .
(NH,OH] =kyus (I)
The value of this constant has been found to be about 2 at 20°. The

ammonium hydroxide, however, is not only in equilibrium with ammonia,
but also with ammonium and hydroxy] ions,

NH,OH — NH,t+0H-,
and the mass-action expression of this equilibrium is
[NH,*] [OH]

[NH,OH]

Moore * found the value of this constant to be about 5X10—. By adding
[NH,0H]

= kpase (II) °

to the left-hand member of equation (I) and its equivalent 1 to the

[NH,OH]
right-hand member, we obtain the equation
[NH;]+[NH,OH] _ pee
vH;|+[NH,OH
and [NH,oH) =#! a tit)

* J. Chem. Soc., 91 (1907), 1379.

AMMONIUM 89

Inserting this value for [NH,OH] in (II) and transferring k’ to the right-hand
member, we have

| [NHAOH=]__ frase _ 7
[NH.OH]+[NH;] k 9”

The value of this constant, K, at 18° is about 1.8107.

The ammonium salts, unlike the free base, are largely ionized. When,
therefore, an excess of OH is added to the solution of an ammonium salt,
it is necessary, in order to establish equilibrium between NH,* and OH™ for
the greater part of the NH,* to be converted into NH,OH and then, to establish
equilibrium between NH,OH and NHs, about two-thirds of the NH,OH is

changed into NH;. By boiling the solution, the NH; is expelled and the
above-mentioned states of equilibrium-are disturbed and, as the final result,
all of the original NH,t becomes converted into NH; gas. Less than 0.2 mg.
of ammonium can be detected by the litmus test when properly carried out.
Care should be taken not to boil the solution so that some of the alkaline liquid
becomes spattered into the nostril or upon the test paper.

Certain complex ammonia derivatives do not always evolve ammonia

aes 2 ; /NEib .
in this test. When pure mercuric amidochloride, Hg , 1s heated

with caustic soda solution, a part of the nitrogen is evolved as ammonia,
but if considerable mercuric salt is present the test is not obtained.
This is because the mercuric amidochloride itself is only slightly soluble
in water and, especially in the presence of an excess of mercuric. com-
pounds, furnishes scarcely any ammonium ions. If some potassium
sulfide is added, however, the mercury is converted into more insoluble
mercuric sulfide and the ammonia test is obtained:

NH,
HeC +K>8+H»0 =HgS+KCl+KOH+NH; f.
Cl

Water itself in some cases causes evolution of ammonia gas. It
decomposes many nitrides, metal amides and cyanamides:

Mg3N2+6HOH =38Mg(OH)2+2NH3 17,
NH2Na+HOH =NaOH+NHs317,.
CaCN2+3H20 = CaCO3+2NHs3 ¢ (at high temperatures)

Calcium
cyanamide

2. Chloroplatinic Acid gives a yellow crystalline precipitate:
H2[PtCle]++2NH4* > (NH4)2PtCle+2H".

This salt may be distinguished from the potassium salt—
(a) by its behavior on ignition; platinum alone is left behind:

(NH4)2[PtCle] =2NH4Cl+2Cle T +Pt;

90 REACTIONS OF THE METALS

(b) by its behavior on treating with strong bases, whereby the smell
of ammonia may be detected:

(NH4)o[PtCle]-+2Na0H = Nas[PtCle]+2H20+2NHz3 Tf.

3. Tartaric Acid produces, as with potassium, a white, crystalline ~
precipitate of ammonium acid tartrate. The addition of a little sodium
acetate, and rubbing the sides of the glass vessel with a stirring-rod, will
hasten the formation of the precipitate: 3

HoC4H406.+NH;t > NH.HC.Hi0,¢+H".

The ammonium acid tartrate, like the corresponding potassium
salt, is soluble in alkalies and mineral acids. It may be distinguished
from the potassium salt by its behavior on ignition, carbon alone
being left behind, and the residue not effervescing with hydrochloric
acid; furthermore, ammonium acid tartrate will give off ammonia on -
Cine heated with caustic soda solution.

4. Sodium Cobaltinitrite gives a yellow pr ecipitate similar to
that produced with potassium. Before testing for potassium with
this reagent, therefore, it is necessary to expel ammonium salts by
evaporating the solution to dryness in a porcelain dish and heating
until no more fumes are evolved.

The above-mentioned reactions are not suitable for the detection of the
very small amounts of ammonia or of ammonium ions that are found in drinking-
water. Insuch cases Nessler’s reagent is used (an alkaline solution of potassium
mercuric iodide). Large amounts of ammonia produce a brown precipitate,

The precipitate is the iodide of the so-called Millon’s base; its structural
Hg
formula is probably 0< Nt -I. The corresponding black nitrate is formed
H

in the test for ammonia with mercurous nitrate paper. . The iodide of Millon’s
base has such a remarkable coloring power that mere traces of the ammonium
ion can be detected by the yellow or brown color imparted to the solution. The
test is obtained with ordinary distilled water. Since ammonia is often present
in water as a result of its contact with decaying organic matter, the test for
ammonia helps to determine whether a water is suitable for drinking. The test
is called the Nessler test, and the alkaline solution of potassium mercuric iodide
is called Nessler’s reagent.

Water free from ammonia should be used in carrying out the Nessler test.
Some sodium carbonate and a little potassium permanganate is added to ordi-
nary distilled water, which is then redistilled, rejecting the first fourth and last
sixth of the distillate; the middle portion is the so-called best water of the
chemical laboratory. For the most accurate work, Nessler’s reagent should

AMMONIUM 91

be prepared with such water (cf. p. 73) and the test should be made in a labora-

tory from which ammonium fumes are absent.

To test a water for the presence of traces of ammonium ions, the apparatus
shown in Fig. 4 may be used. First of all, the apparatus itself must be freed
from all traces of ammonium salt. To accomplish this, place about 50 cc. of
water in the retort, add 1 ce. of a boiled, saturated solution of sodium carbonate
and distill with the neck of the retort introduced well into the condenser tube.
It is advisable not to use a rubber connection between the retort and the con-
denser; the condensed water serves to make a sufficiently tight connection.

Continue distilling until 50 c¢. of the distillate placed in a white glass
graduate, or in a so-called Nessler tube, and treated with 1 cc. of Nessler reagent,

YA
) Sarace

| | | |
H
EI

Fig. 4.

will show no sign of color after standing five minutes. The apparatus is then
ready for the test.

Empty the retort and refill it with 500 cc. of the water to be tested, add 1 cc.
of the saturated sodium carbonate solution, distill and collect the first 50 cc.
of distillate. If as much as 0.2 mg. of ammonium is present, a distinct precipi-
tate forms on adding 1 cc. of Nessler’s solution and stirring; a pronounced
yellow color is obtained with much smaller quantities. If mere traces of
ammonium are present, the yellow color appears only on standing. By com-
paring the depth of color with that similarly produced with known quantities
of ammonium chloride, a very close estimate of the exact quantity of ammo-
nium present can be made,

REACTIONS IN THE DRY WAY

All ammonium salts are relatively unstable compounds, the degree
of stability depending, in general, upon the strength of the acid which
is combined with the ammonium. The carbonate decomposes appre-
ciably at ordinary temperatures and when exposed to the air gradually

92 REACTIONS OF THE METALS

. disappears as ammonia, carbon dioxide and water. Heating in a closed
tube causes the decomposition of all ammonium salts and either
ammonia or some other volatile nitrogen compound escapes.

If the acid is volatile at the decomposition temperature, the whole
salt is volatilized, often without first melting, and when the vapors
of ammonia and acid are cooled the solid again forms. This explains
the ammonium fumes that result when ammonium salts are expelled
from a solid residue obtained by evaporation of a solution, and it
explains the sublimate formed when ammonium chloride is heated in
a closed tube. The acids which form salts that are not volatilized are
boric, phosphoric, chromic, molybdic, tungstic and vanadic acids.

It must be remembered, however, that ammonia with its negative
valence of three contains nitrogen in its lowest state of oxidation. When
the decomposition of the original ammonium salt takes place, therefore,
there is often an oxidation of the nitrogen. Thus the decomposition of
the nitrate results in the formation of nitrous oxide, N2O, and the
decomposition of the nitrite, sulfate and dichromate yields nitrogen gas.

The closed tube reactions of typical ammonium compounds may be
expressed by the following equations:

NH,Cl = NH3+HCl,
NHiNO;=2H20+N;0,
NHiNO2=2H20-+Np,
3(NH1)2804 = No-+4NH3+6H20+3802,
(NH) 2C204=2NH3+H20+CO+COr.

Toward the last some dicyanogen, (CN)s, is formed from the oxalate.

NaNH4HPO, : 4H2O => NH3 +5H»eO +NaPOs,
2(NH1)3PO,=6NH3+2H20+2HPO3,
(NH4)2Cr207 =4H20+Ne2+CreOsz.

In this last reaction, the chromic oxide remains as a voluminous mass
looking something like green tea. A realistic voleano effect can be
obtained by making a mound of ammonium dichromate around a piece
of paper and then setting fire to the latter.

Ammonium salts do not impart a characteristic color to the flame;
the border of the flame is tinged slightly greenish.

MAGNESIUM 93

MAGNESIUM, Mg. At. Wt. 24.32
Sp. Gr. =1.75. M. Pt. =632.6° C.

Occurrence—Magnesium compounds are found very abundantly
in nature. The most important minerals are magnesite, MgCOs,
rhombohedral, isomorphous with calcite; dolomite, (Ca,Mg)CO3;
brucite, Mg(OH)2, rhombohedral; carnallite, KMgCls+6H20, ortho-
rhombic; kieserite, MgSO4+HeO, monoclinic; epsomite, MgSO,
+7H20, orthorhombic; spinel, MgAloOx4, isometric, isomorphous with
magnetite, Fe3O04, and with chromite, FeCreO4. Magnesium also occurs
in a great many silicates. Thus almost all the minerals of the olivine
group contain more or less magnesium. To this group belong forsterite,
Mg2S8i04; monticellite, CaMgSi0O4; and olivine, FeMgSiO«z; all ortho-
rhombic. An important decomposition product of the olivine minerals
is serpentine, Mg3H4Si2O9. Almost all the minerals of the pyroxene-
amphibole group, which are all related to orthorhombic enstatite,
MgSi03, contain magnesium: augite, MgAloSiOg; hornblende, an
isomorphous mixture of Mg3CaSi4Oi2 and 2(MgAl2Si0¢) ; and tremolite,
CaMg3SisO12, all three forming monoclinic crystals. Asbestos is a
variety of tremolite with very fine fibers. Meerschaum is a magnesium
silicate of the composition HsMgeSisOi0, and is quite similar to tale,
H2Mg3SisO12, sometimes called steatite. Magnesium also occurs in
the vegetable kingdom, being an essential constituent of the complex
organic compound chlorophyll.

Properties of Magnesium.—Magnesium is a silver-white metal.
It decomposes water very slowly, forming an oxide, MgO, which
is only slightly soluble in water, forming magnesium hydroxide; the

small quantity that dissolves is largely ionized, so that the solution has

a faint alkaline reaction. Magnesium reacts directly with nitrogen at
300° C., forming magnesium nitride (Mg3N2), which is readily decom-
posed by water, forming magnesium hydroxide and ammonia:

Mg3N2-++6HOH =3Mg(OH)2+2NHs3 fT .

The salts of magnesium are almost all colorless and soluble in water.
The solubility products of magnesium hydroxide, carbonate, phosphate,
arsenate and arsenite are so small that these substances may be regarded
as insoluble. The sulfide, which can be prepared only in the dry way,
is completely decomposed by water into hydroxide and hydrogen sul-
fide (hydrolysis). If an aqueous solution of magnesium chloride be
evaporated to dryness on the water-bath, there is no hydrolytic decom-
position, the residual salt, MgCl2+6H20O, dissolves in water, forming a
perfectly clear solution. On heating the chloride containing the water

94 REACTIONS OF THE METALS
fx

of crystallization to 106° and higher, however, a considerable amount
of hydrochloric acid fumes escape and a basic magnesium chloride
insoluble in water is left behind:

Cl
MENS
2MgCle +He2O —_ O +2HCIl.
Me
Nol

When a saturated solution of magnesium chloride is mixed with
magnesium oxide, the mixture soon solidifies, forming a mass hard
as stone, known as magnesia cement, consisting of basic magnesium
chloride.

REACTIONS IN THE WET WAY

1. Strong Bases, such as the soluble hydroxides of sodium, potassium
and barium, precipitate white, gelatinous magnesium hydroxide, and the
precipitation is practically complete in the absence of ammonium salts
or if the ammonium salt is all decomposed by boiling with an excess of
the strong base (cf. pp. 19, 46).

Mgt ++20H- — Mg(OH)2

The solubility product of magnesium hydroxide (cf. p. 22) is about 3.4107"
at the laboratory temperature. The saturated solution of magnesium hydrox-
ide in pure water contains about 0.0002 mole or 0.012 gm. Mg(OH), per liter.
In the presence of an excess of OH, the solubility of the magnesium hydroxide
is much less, as a result of the common ion effect (cf. p. 45) and it is possible,
by keeping the volume of the solution small and using a slight excess of the
reagent, to leave less than 1 mg. of magnesium in solution.

The precipitation of magnesium hydroxide by means of the slightly ionized
ammonium hydroxide can never be made complete, and if the solution already
contains ammonium ions in sufficient excess, no precipitation of magnesium
hydroxide takes place. Moreover, if a precipitate of magnesium hydroxide is
boiled with a solution of an ammonium salt, such as ammonium chloride, the
precipitate dissolves. This behavior is due to the fact that the ionization of
ammonium hydroxide (cf. p. 19 and p. 88) is repressed to such an extent,
as a result of the common ion effect, that not enough OH~ ions are present at
any one time to satisfy the solubility product of magnesium hydroxide, and
even the OH~ from the Mg(OH),. must be in equilibrium with the NH,* ions
from NH,Cl.

Formerly, the non-precipitation of magnesium by ammonium hydroxide
was explained by the assumption that complex salts such as NH.{[MgCl,] or
(NH,).[MgCl,] were formed, but this explanation has proved untenable.*

* Cf. Loven, Z. anorg. Chem., 11 (1896), 404; TreapweE.., ibid., 37 (1903), 326
and Herz, ibid., 38 (1903), 138.

MAGNESIUM 95

_ 2. Mercuric Oxide heated with solid magnesium chloride converts
the latter into magnesium oxide which does not dissolve appreciably in
water. The mercuric chloride formed and the excess of mercuric oxide
are volatilized.

3. Ammonium Carbonate precipitates, in the absence of other
ammonium salts, a basic salt (usually only on boiling or after long
standing). The composition of the precipitated salt varies with the
temperature and the concentration of the solution, the following salt
being often obtained:

4Megtt +4(NH4)2CO3 +H20 — Mg, (CO3)3(OH)2+CO2T +8NH4,",

The addition of an excess of ammonium carbonate reagent (p. 72)
and an equal volume of 95 per cent alcohol causes the complete pre-
cipitation of magnesium as MgCO3- (NH4)2CO3-4H20 from a cold, con-
centrated solution of magnesium salt.

The magnesium ammonium carbonate is fairly soluble in water
and the solubility increases rapidly with rise of temperature. Thus
no precipitate is obtained upon the addition of ammonium carbonate
to a hot dilute solution of magnesium salt containing ammonium
chloride and no alcohol. (Note difference from barium, strontium
and calcium.) |

' 4. Sodium Phosphate is the characteristic reagent for magnesium.
It produces in solutions containing ammonium chloride, and in the
presence of ammonia, a white crystalline precipitate (orthorhombic,
hemimorphous) of magnesium-ammonium phosphate, *

Mgt Fo NH4t+ POw— MgNH,PO,.

From very dilute solutions the precipitate separates only after
standing some time, owing to the tendency to form supersaturated solu-
tions. Rubbing the sides of the beaker with a glass rod hastens the
formation of the precipitate.

The table on p. 22 gives the solubility product of magnesium ammonium
phosphate as 2.5X107~; about 0.0086 mg. of the salt dissolves in a liter of
water at the laboratory temperature. The solubility in’water is increased by

its tendency to undergo hydrolysis.
MgNH,P0.,+HOH @ Mgt *++HPO~+NH.OH. (I)

The tendency to undergo hydrolysis increases rapidly with rise in temperature
(cf. p. 51). A similar decomposition is caused by hydrogen ions alone,

MegNH,PO,+Ht @ Met++HPOC+NH.t, (II)

* Magnesium ammonium phosphate crystallizes with six molecules of water.

96 REACTIONS OF THE METALS

which lessen the tendency for HPO, to ionize; the precipitate, therefore, dis-
solves readily in the presence of any acid which is ionized more than HPO,-, even
acetic acid (cf. p. 10). The presence of ammonium hydroxide prevents the
hydrolysis, in accordance with the mass-action law. For this reason an excess
of ammonia solution is usually added. Ammonium chloride, by virtue of the
common ion effect, lessens the quantity of Mgt* ions required to reach the
solubility product and causes reaction (II) to take place in the direction right
to left; but, on the other hand, it should favor reaction (I) somewhat, because
it represses the ionization of ammonium hydroxide. As a matter of fact
ammonium salts usually retard the formation of the precipitate, but do not
eventually make it more soluble if an excess of ammonia is present.

Neubauer * has shown that the conditions are still more complicated on ac-
count of the tendency for gelatinous Mg;(PO,). and crystalline Mg(NH,)4(PO,)2
toform. Tertiary magnesium phosphate, Mg;(PO,)s, is formed in cold, strongly
ammoniacal solutions containing but little ammonium salts. The mono-
magnesium-tetrammonium phosphate is formed in neutral or slightly alkaline
solutions containing considerable ammonium salts. B. Schmitz? has shown
that beautifully crystalline precipitates can be obtained in the presence of
ammonium salts by adding sodium or ammonium phosphate to the boiling,
acid solution of the magnesium salt. Then, on adding one-third the solution’s
volume of 6 N ammonia and allowing the solution to cool, complete precipi-
tation, as MgNH,PO,-6H,0, takes place.

REACTIONS IN THE DRY WAY

All magnesium salts are more or less changed on heating in the
air, leaving behind the oxide or an insoluble basic salt. On char-
coal with sodium carbonate before the blowpipe, magnesium com-
pounds are changed to white magnesium oxide, which is strongly
luminous when hot. Calcium, strontium, and aluminium compounds
behave the same way. The magnesium salts are nonvolatile, do not
color the flame, and give no flame spectrum, but do give a characteristic
spark spectrum.

Detection and Separation of Magnesium and the Alkalies in the Pres-
ence of One Another

All of the typical reactions that have been described for magnesium
can be used for separating magnesium from the alkalies. The test
for ammonium is always carried out with some of the original substance
or solution. |

Ordinarily in qualitative analysis, it is customary either to test for
magnesium in one portion of the solution by means of sodium phosphate
and for sodium and potassium in another portion, or to precipitate the

* Z. angew. Chem. (1896), 439. Cf. Goocn and Austin, Z. anorg. Chem., 20, 121.
+ Z. anal. Chem. (1906), 512. . Cf. JonaEnsEN, ibid. (1906), 278.

MAGNESIUM 97
magnesium as carbonate and test for sodium and potassium in the
filtrate. In this last method the magnesium can be precipitated with
the alkaline earths. If this is done, the test for magnesium is not made
in the analysis of this group.

In all the principal systematic methods given in this book for the
examination of the basic constituents, a very brief tabular outline
will be given first with reference by numbers to the detailed direc-
tions that immediately follow.

TaBLE I.—ANALYSIS OF Group V. Mertuop A.

NH;+. Test a|Mgt*t. Test a| Evaporate to dryness. Expel NH, salts.
portion of the} part ofthe fil- Dissolve in water, add Ba(OH)2 and reject
original sub-| trate from| _ precipitate. Add. HCl, NH,OH and
stance for| Groupll/V, (NH4)2CO; and also reject this precipitate.
NH; by boil-| Method A, for Evaporate, expel NH, salts, dissolve in water,
ting with| Mg with| filter, and add HCl0O,4. Evaporate, add
NaOH. (1) NH,OH_ and alcohol and filter. (8).

N. a2H P O14. (2)

NazH.Sb2O7. (5)

PROCEDURE

1. Test for Ammonium. Place a little of the original substance (correspond-
ing to about 0.25 gm. of solid) in a test-tube, add about 2 cc. of 6 N sodium
hydroxide, heat nearly to boiling and hold a piece of red litmus paper, wrapped
around the end of a stirring rod, in the escaping vapors. Take care not to allow
any of the caustic alkali to come in contact with the litmus either by spattering
or by allowing the paper to touch the sides of the test-tube. A good idea of the
quantity of ammonium present can be estimated by the odor. If it is desired
to test for traces of ammonium, carry out the Nessler test as described on p. 91.

2. Test for Magnesiuin. Dissolve the substance in as little water as possible,
or, if a solution, evaporate to dryness, moisten with 6 N hydrochloric acid,
heat gently, dissolve in a little water and filter if necessary. The addition of the
acid is not usually necessary, but sometimes, when calcium has previously been
removed as oxalate, a difficultly soluble oxalate of magnesium and ammonium
is formed which is best dissolved by treatment with acid before adding water.
If the solution to be tested contains considerable ammonium salt, it is usually
best to expel it by igniting the residue obtained by evaporation; the treatment
of the residue with hydrochloric acid is then absolutely necessary, in order to
decompose basic magnesium salts which are formed during the ignition.

If ammonium chloride is not already present, add half the solution’s volume
of normal ammonium chloride and enough ammonia to make the solution dis-

98 REACTIONS OF THE METALS

tinctly ammoniacal. If a precipitate of magnesium hydroxide is formed, add
more ammonium chloride to dissolve it. A flocculent precipitate produced at
this point may be aluminium hydroxide or silica. Such a precipitate should
be filtered off and discarded. To the clear ammoniacal solution add a little
sodium phosphate solution and rub, with a rounded stirring rod, the inside
walls of the vessel containing the solution. If as much as a few tenths of a
milligram of magnesium is present per 100 cc., a crystalline precipitate of
magnesium ammonium phosphate should form within a few minutes. In case
no precipitate is noticeable, set the beaker aside and allow it to stand overnight.
Traces of magnesium phosphate will usually form on the sides of the beaker
where it was scratched by the stirring rod.

If the precipitate does not appear distinctly crystalline, it may contain
aluminium phosphate. In such cases filter off the precipitate and redissolve
it in 6 N acetic acid, which will not dissolve aluminium phosphate, but
readily dissolves magnesium ammonium phosphate. Heat the solution nearly
to boiling, add a liberal excess of ammonia and allow to cool; a crystalline pre-
cipitate should be obtained by this treatment.

3. Test for Sodium and Potassium. Before testing for sodium and potassium
it is best to remove the magnesium. Evaporate the solution to dryness in a
porcelain or platinum dish and gently ignite the residue to expel ammonium
salts. Dissolve the residue in a little water and, without paying any atten-
tion to any residue of basic magnesium salt, add barium hydroxide solu-
tion until strongly alkaline. Heat to boiling and filter off the precipitate, which
may contain magnesium hydroxide, barium sulfate, if any sulfate ions were
present, and barium carbonate from the contact of the barium hydroxide
solution with the air. These operations should never be carried out in glass
dishes because of the danger of obtaining alkali from the glass. Make the fil-
trate from the magnesium precipitate barely acid with hydrochloric acid and
remove all barium ions by the addition of ammonia and ammonium earbon-
ate. Filter off the precipitated barium carbonate, evaporate the filtrate to
dryness and expel the ammonium salts by careful ignition. Dissolve the resi-
due in a little water and make sure that all the barium was removed by add-
ing a little more ammonium carbonate. Filter if necessary and again evapo-
rate to dryness. Expel the ammonium salt by ignition, cool and wash down
the sides of the dish with about 5 cc. of water to dissolve any ammonium salt
that may possibly be left there. Evaporate to dryness again and expel the last
traces of ammonium salts by igniting at a dull red heat until no more fumes
are evolved. Dissolve the residue in a little water and filter the resulting
solution through a small filter. A black residue is often due to carbonization
of pyridine bases, which are commonly present to a slight extent in ammonia
solutions.

Evaporate the solution just to dryness, add 5 to 15 ce. of 2 N perchloric
acid and evaporate carefully, by keeping the dish in constant motion over a
free flame, until dense fumes of perchloric acid are evolved. Cool completely
and add 20 ce. of alcohol. .

The perchloric acid solution must not be heated or concentrated by evaporation
after the alcohol has been added or a dangerous explosion is likely to result.

Stir the solution and press down gently with a stirring rod upon any crystals
that may be present. If after stirring a few minutes a residue of potassium
perchlorate remains, add 3 cc.-more of the perchloric acid reagent. Filter
through a dry filter-paper and wash the precipitate with alcohol. vem

,
{
F

MAGNESIUM 99

4, If the previous manipulation has been faulty, the precipitate obtained
with perchloric acid-may consist of ammonium perchlorate, and it is advisable,
therefore, to confirm the potassium test. Dissolve the precipitate on the filter
with as little hot water as possible, using not over 10 cc. at the most. Add
the water in small portions around the top of the filter paper and pass the first
filtrate through the filter a second time. Add 8 to 5 ce. of sodium nitrite
reagent and 2 to 3 cc. of 6 N acetic acid and boil gently for about five minutes.
The boiling with nitrous acid serves to decompose any ammonium salt present;
it will cause the decomposition of as much as 30 mgs. of ammonium ions.

NH,++NO.-— N, t +2H.0.

Cool the solution and add a little sodium cobaltic nitrite reagent. A yellow
precipitate of K,NaCo(NO2). will form within ten minutes if as much as 0.3
m. of potassium is present.

5. Test for Sodium. Pour the alcoholic filtrate from the perchlorate pre-
cipitate into a small Erlenmeyer flask, place the flask in cold water atid saturate
with dry hydrogen chloride gas. This gas may be prepared by dropping con-
centrated sulfuric acid into a flask containing common salt and concentrated
hydrochloric acid and passing the escaping gas through concentrated sulfuric
acid. In the presence of alcohol the hydrogen chloride will precipitate a
little as 1 mg. of sodium ions as sodium chloride. Filter off the precipitated
sodium chloride and wash it with a little alcohol. Dissolve it in a very little
water, evaporate the solution to dryness, dissolve in 1 cc. of water and add
twice as much potassium pyroantimonate reagent. A crystalline precipitate

of sodium pyroantimonate can be obtained with as little as 1 mg. of sodium ions

if the solution is allowed to stand overnight. Traces of many other elements,
such as the alkaline earths, magnesium and aluminium also give precipitates
with this reagent, but such precipitates are flocculent.

The final precipitates of potassium and sodium should always be submitted
to the flame test in all cases of doubtful precipitation.

TaBie II.—ANAtysis oF Group V. Mernop B.

NH.+ Testthe|_.
Sri Nets Filtrate from Group IV, Method A, may

et Method A. contain Mg++, Kt, Nat, NHy+. Concentrate to 10 cc.,
p. 97. ’| add 15 cc. (NH4)2COs and 15 cc. CxHyOH. (1)

Precipitate:] Filtrate: K+, Nat, NH,+. Remove sul-
MgCoOs3:(NH,4)s fate by a BaCh, remove Bat+t with
CO;:4H:0. Dis-| (NH,),CO;. Expel NH, salts and
solve in 6-normal add HCIO,. (3)

H»SO, and add an

equal volume of ae
C.H;OH. Filter, | Precipitate: KC1O,, | Filtrate, Nat. Test

add NH,OH and| Examine as in| for Na as in
NasH PO; to pre- Method A, Dp. 98. Method A, p. 98.
cipitate Mg

a Law

Na ee ee ee

100 REACTIONS OF THE METALS

PROCEDURE*

1. Concentrate the solution to-a volume of about 10 ec. and filter if necessary.
If ammonium salts are deposited it is best to remove them first as in Method A.
Add to the concentrated solution 15 cc. of ammonium carbonate reagent and
15 cc. of alcohol, stir well, and allow the mixture to stand half an hour or longer,
Filter off the precipitate of MgCO;-(NH,)2.CO;-4H.0.

2. To confirm the magnesium test, dissolve the precipitate, which may
contain alkaline-earth carbonate, in a little 6 N sulfuric acid and add an —
equal volume of alcohol. Any barium, strontium or calcium will be converted
. into insoluble sulfate by this treatment. Filter the solution if necessary and
test for magnesium with sodium phosphate solution in the usual way.

3. Test for Potassium and Sodium. Evaporate the filtrate from the precipi-
tate obtained by treatment with ammonium carbonate, expel ammonium salts
and if a sulfate is present, heat to boiling and add hot barium chloride solution —
until there is no further precipitation. To the filtrate add (NH4)sCO; solution,
heat to boiling and filter off the BaCO;. Evaporate the filtrate to dryness in a
porcelain dish and heat the residue until no more white fumes are evolved.
Cool completely, add 5 cc. of water, washing down the sides of the dish with it,
filter off the carbonaceous residue, evaporate to dryness and again heat to make
sure that all the ammonium salts are expelled. Cool, add HCIO,, and treat as
directed in Method A, p. 98. If no sulfate is present it is unnecessary to add
BaCl, and (NH,)2COs;, otherwise the procedure is the same.

The separation of the potassium and sodium can be accomplished
by treatment with chloroplatinic acid instead of perchloric acid, but
such a method is not suitable in ordinary qualitative analysis on account
of the expense involved. The treatment is practically the same except
that instead of evaporating till fumes of perchloric acid are evolved,
the solution is evaporated just to dryness on the water-bath.

*Cf. Scuarrcorr, Ann. Phys., 194 (1858), 482; Goocu and Eppy, Z. anorg.
Chem. (1908), 427; A. A. Noyes, A Course of Instruction in the Qualitative
Analysis of Inorganic Substances (1914).

GROUP IV. ALKALINE EARTHS

CALCIUM, STRONTIUM, BARIUM

GENERAL CHARACTERISTIC REACTIONS

The metals of the alkaline-earth group are bivalent and heavier
than water, which they decompose slowly at ordinary temperatures,
with evolution of hydrogen and formation of difficultly soluble hydrox-
ides of strongly alkaline reaction (cf. p. 41). The salts are mostly
colorless and very slightly soluble in water. The halogen compounds,
nitrates, nitrites, and acetates are soluble in water. The carbonates are
insoluble in water and are decomposed on ignition into carbon dioxide,
and white, infusible, strongly luminous metallic oxide:

CaCO3 = Ca0+C02}.

Strontium carbonate is less readily decomposed than calcium carbonate,
and barium carbonate loses its carbon dioxide only when heated to a
white heat; its oxide is not very luminous.

The sulfates and oxalates are very difficultly soluble. The
sulfate of barium is the most insoluble sulfate and calcium sulfate the
most soluble; of the oxalates, the calcium salt is the most insoluble
(ef. p. 21). The solubility of the strontium salt is always midway
between that of the corresponding calcium and barium salt, for the
atomic weight of strontium, of which the solubility is a function, lies
midway between the atomic weights of barium and calcium. The halo-

gen salts are volatile and impart a characteristic color to the flame.

The metals of this group form oxides of the general type RO, and
peroxides corresponding to the formula ROz. The latter, on treatment
with acids, give hydrogen peroxide and salts corresponding to the oxide

RO:
ROe+2HCl = H202+RCle.

Magnesium is more closely related to this group than to the alkali
metals. It can be precipitated with this group if the group precipitant,
ammonium carbonate, is added to the concentrated solution together
with an equal volume of alcohol.

101

102 REACTIONS OF THE METALS |

CALCIUM, Ca. At. Wt. 40.07
Sp. Gr. 1.58. M. Pt. 810°C.

Occurrence.—Calcium is widely distributed in nature. It is found
in enormous deposits in all stratified formations as carbonate (limestone,
marble, chalk), often rich in petrification. The carbonate, CaCOs,
is dimorphous, crystallizing in rhombohedrons as calcite and in the
orthorhombic system as aragonite. Calcium also occurs in large
masses as sulphate, partly as monoclinic crystallizing gypsum,
CaSO.4-2H20, and partly as anhydrite, CaSO4, which crystallizes in the
orthorhombic system. Calcium also occurs as fluorite, CaFe, which
crystallizes in the isometric system, with perfect octahedral cleavage;

Cl
as apatite, 3Cag(PO.)2: Cat , which belongs to the hexagonal system;
KF .

and, finally, in innumerable silicates, such as the monoclinic wollaston-
ite, CaSiO3, and the triclinic anorthite, CaAlsSis0s. The calcium
minerals are the principal representatives of several important
isomorphous groups:

Calcite Group, Rhombohedral. Aragonite Group, Orthorhombie.

Calcite, CaCOsz Aragonite, CaCOz
Magnesite, MgCOs Strontianite, SrCOz3

: Ca Witherite, BaCOsg
pepomite, Mg as Cerussite, PbCOg3
Siderite, FeCO3

Smithsonite, ZnCOz3
Rhodochrosite, MnCOg

Anhydrite Group, Orthorhombic. Apatite Group, Hexagonal.

Anhydrite, CaSO4 Apatite, 3Caz(PO4)2+Ca(ClF)
Celestite, SrSO4 Pyromorphite, 3Pb3(PO4)2+PbCle
Barite, BaSO, Mimetite, 3Pb3(AsO4)2+PbCle
Anglesite, © PbSOs Vanadinite, 3Pb3(VO4)2+PbCle

REACTIONS IN THE WET WAY

1. Ammonia, in case it is free from carbonate, produces. no pre-
cipitate with calcium salts; on standing in the air, however, carbonic
acid is absorbed and a turbidity of calcium carbonate results.

2. Ammonium Carbonate, or any other soluble, normal carbonate,
precipitates white calcium carbonate; the precipitate is voluminous

CALCIUM | 103

and flocculent when it first forms, but soon becomes crystalline, par-
ticularly when in contact with boiling water.

Cat++CO37* — CaCO3. (1)

‘The precipitate is noticeably soluble in an aqueous solution of the

ammonium salt of any strong acid:
CaCO3+2NH4t — Catt++2NH3 T 41,0 +00: 1 a (IT)

When ammonium carbonate is the precipitant, equation (II) is essen-
tially the cpposite to equation (I) and the mass action law shows how
the reaction can be made to go in either direction. An excess of am-
monium carbonate will favor the progress of equation (1) and boiling
with a large quantity of an ammonium salt such as ammonium chloride
will cause equation (II) to go to completion.

Ammonium carbonate is an unstable substance (cf. p. 51). Commercial
ammonium carbonate, often called ammonium sesquicarbonate, is a mixture
of approximately equivalent quantities of ammonium bicarbonate, NH,HCOs,
and ammonium carbamate, NH.CO.NH2; the latter salt corresponds to normal
ammonium carbonate less one molecule of water. Calcium carbamate is quite
soluble in water, but by contact with water at 60° it becomes changed to insolu-
ble calcium carbonate. The ammonium carbonate reagent is prepared with
6 N ammonia instead of water; this prevents the hydrolysis of the salt and
changes the bicarbonate to ammonium carbonate.

Calcium bicarbonate is soluble in water, so that any acid which is
dissociated to a greater extent than HCOs3 will exert a solvent effect
upon calcium carbonate.

CaCO; +H Fi Cat Fi. HCO3.

Thus acetic acid dissolves calcium carbonate readily. Boiling the solu-
tion favors the progress of the reaction, as the HCO37 is also in equi-
librium with H* and H2CO3 and the latter with H2O and COs:

HCO;"+H* — HeCO3 — H20+COd2,

and the carbon dioxide can be expelled completely at 100° (ef. p. 15).

The precipitation of calcium carbonate from boiling, dilute solutions
containing ammonium salts is always more or less incomplete, for
the reasons that have just been given, but from cold, concentrated
solutions containing considerable alcohol the precipitation is practically
complete.

3. Ammonium Oxalate produces in neutral or alkaline solutions
a precipitate of calcium oxalate, which when formed from cold solutions

104 i: REACTIONS OF THE METALS

is composed of extremely fine crystals, hard to filter, while from hot
solutions larger crystals are formed:

C2047 +Catt = CaC20,.

Calcium oxalate is practically insoluble in water and aectin acid, but
dissolves readily in mineral acids:

CaC.04+2Ht 2 Catt++HoCs0u.

Calcium oxalate, unlike calcium carbonate, does not dissolve in
acetic acid for the following reasons: (1) The solubility product of
calcium oxalate is about 3.8x10-® while that of calcium carbonate is
about 1.7X10-® (p. 21) which shows that the oxalate is somewhat
less soluble in water than the carbonate. (2) The ionization constant
of HCO3° is 0.0638 while that of HC2047 is 0.045 and that of acetic acid
is 0.0418; this reason is sufficient to explain why acetic acid has little
solvent effect upon calcium oxalate. To dissolve calcium oxalate readily
it is necessary to use an acid strong enough to repress the ionization of
the first hydrogen of oxalic acid, for which the ionization constant is
0.038. (38) The progress of the reaction cannot be aided, as in the case
of the carbonate, by the loss of a volatile constituent.

Ammonia precipitates from such a solution calcium oxalate again,
the excess of hydrogen ions, as well as, the oxalic acid which was formed,
being neutralized.

Calcium oxalate on being boiled with sodium carbonate solution
is easily changed to carbonate:

CaC204+C0O37 @ CaCO3+ C2047.

This reaction takes place in the direction left to right when an excess
of COs” ions are present in spite of the fact that calcium oxalate is
somewhat less soluble than calcium carbonate. This is in accordance
with the mass action law (p. 13). An excess of C2047 ions will make
the reaction take place in the direction right to left. |

4, Sulfuric Acid produces a precipitate only in concentrated
solutions:

CaCle +He2SO4 = 2HCl+CaSO..

One hundred cc. water disolves 0.214 gm. CaSO4-2H2O at 40°,
but much less than 1 mg. will dissolve in the same quantity of alcohol.
In the presence of a slight excess of sulfate ions, calcium sulfate is less
soluble, but the behavior of calcium sulfate toward hydrogen ions is
similar to that of calcium carbonate. Calcium acid sulfate is much
more soluble than normal calcium sulfate, and any acid which is capa-

sa eee
; 7 -

CALCIUM , 106

ble of repressing the ionization of HSO,7 will exert a solvent effect
upon CaSO4. Sulfuric acid itself, or any other strong acid, can exert
this solvent effect.

~ Calcium sulfate is also soluble in concentrated ammonium sulfate,

owing to the formation of a complex anion:

CaSO4+(NH4)2804 => (NH4)2[Ca(SO4)o].

5. Calcium Sulfate solution produces no precipitation with calcium
salts. (Note difference from strontium and barium.)

6. Sodium Phosphate (NazHPOs) produces in teaibcal solutions
a white, flocculent precipitate of secondary calcium phosphate:

Catt++HPO,- — CaHPOx,.

If ammonia is added to the solution at the same time, tertiary
calcium phosphate will be precipitated:

HPO, +OH — He0+P0.4-
Cat t++2P04* => Cazg(POx)o.

Both of these phosphates of calcium are dissolved by hydrogen ions.
A glance at the table on p. 10 shows that even acetic acid will
repress the ionization of H2PO4" to an extent such that it will exert a
slight solvent action upon CaHPO, and upon Cag(POxs)o. Acetic
acid in the presence of a soluble acetate, however, will not dissolve
Caz(PO4)2 to any extent. From a solution obtained by dissolving cal-
cium phosphate in acid, ammonia always precipitates the tertiary salt.

7. Chromates of the Alkalies do not precipitate calcium salts
from dilute solution. (Note difference from barium and strontium.)

8. Absolute Alcohol, or a mixture of equal parts of absolute alcohol
and ether, dissolves both the nitrate and chloride of calcium.

All deliquescent salts, with the exception of potassium carbonate,
dissolve in absolute alcohol. All other salts are, in general, insoluble,
or very difficultly soluble, in absolute alcohol. An exception to this
rule is found in the case of mercuric chloride, which is not deliquescent,
and is much more readily soluble in alcohol than it is in water.

9. Water decomposes the carbide, phosphide, and nitride of calcium
at the ordinary temperature, as follows:

(a) The carbide:

CaC2+2HOH = Ca(OH) 2-+CoHp 1.

Acetylene
Acetylene is evolved by the reaction, a gas with a peculiar smell.*

* Pure acetylene is odorless. Almost all calcium carbide contains a little ecal-
cium phosphide, which evolves phosphine on treatment with water.

106 REACTIONS OF THE METALS

If this gas is conducted into an ammoniacal copper solution, it rapidly

produces a red precipitate of copper acetylide. The latter compound

is harmless while it is moist; but in the dry state it can be readily

exploded by a blow, by rubbing, or by simply warming.
(b) The phosphide:

- Ca3P2+6HOH =3Ca(OH)2+2PH3 1.

The garlic-smelling phosphine gas which is evolved is spontaneously
combustible, because it always contains a small amount of the spon-
taneously combustible liquid, phosphuretted hydrogen (P2H4).

CasN2+6HOH =3Ca(OH)2+2NH Tf 3.

REACTIONS IN THE DRY WAY

Calcium compounds, on being heated with sodium carbonate
before the blowpipe, are changed to the white infusible oxide, which
glows brightly when hot.

The volatile calcium compounds color the non-luminous gas flame
brick-red.

Flame Spectrum.—Orange-yellow double line (620.3 uy, 618.2 pp)
and a yellowish-green one (554.4 wy, 551.8 wu); both of these lines
belong to calcium oxide. If calcium chloride wet with hydrochloric acid
is placed in the flame, a number of other lines are seen; in the orange-
yellow 646.6 uu, 606.9 uu, 604.5 yu and 593.4 wy, in the yellow 581.7 uy
and 572.0 uu, in the violet, usually very hard to see, 422.7 yy (see
spectroscopic chart, Frontispiece),

STRONTIUM, Sr. At. Wt. 87.63
Sp. Gr. 2.5. M. Pt.>Ca. and<Ba.

Occurrence.—Strontium occurs quite commonly with calcium,
but usually in much smaller amounts. There are only a few true
strontium minerals. The most important of these are: Strontianite,
SrCOz, orthorhombic, isomorphous with aragonite; and celestite,
_ SrSO4, orthorhombic, isomorphous with barite.

REACTIONS IN THE WET WAY

1. Ammonia: same as with calcium.

2. Ammonium Carbonate: same as with calcium.

3. Ammonium Oxalate: same as with calcium; but the stron-
tium oxalate is somewhat soluble in acetic acid.

STRONTIUM 107

4, Dilute Sulfuric Acid produces a white precipitate of -stron-
tium sulfate:
SrCle+-H2SO4 = 2HCl+SrSOx.

Strontium sulfate is much less soluble in water than calcium .
sulfate (6900 parts of water at ordinary temperatures dissolve 1 part

F $rSO.z), but much more soluble than barium sulfate. It is soluble in

boiling hydrochloric acid, and insoluble in ammonium sulfate. By
boiling with a solution of ammonium or alkali carbonate solution, the
strontium sulfate is changed to carbonate:

$rS0O.1+COs — SrCO3+S0.".

5. Calcium Sulfate solution produces in neutral or weakly acid
solutions, after some time, a precipitate of strontium sulfate:

Srt *+-CaSO4 =SrSO4+Cat?.

-6. Chromates of the Alkalies produce in dilute solutions no pre-
cipitate (thus differing from barium); but from concentrated solutions
strontium chromate is precipitated. It is much less soluble in alcohol;
100 ce. of alcohol, 53 per cent by volume, will dissolve 0.002 gm. SrCrO4
and 0.088 gm. of CaCrO4 at room temperature; 100 cc. of 29 per cent
alcohol will dissolve 0.1382 gm. SrCrO4 and 1.22 gms. CaCrO4. The
precipitate is quite soluble in acetic acid.

7. Absolute Alcohol. The nitrate is not deliquescent, and does not
dissolve in absolute alcohol. Strontium chloride is slightly deliquescent;
the anhydrous salt dissolves scarcely at all in absolute alcohol; but,
on the other hand, 1 gm. of SrClz-+6H2O dissolves in 116.4 gms. of cold
alcohol and 262 gms. of boiling absolute alcohol. 100 cc. of 66 per
cent alcohol dissolves about 50 gms. of SrCle-6H20.

REACTIONS IN THE DRY WAY

Heated on charcoal before the blowpipe the strontium compounds
behave similarly to the calcium compounds.

The volatile strontium salts color the non-luminous gas flame
carmine red. |

Flame Spectrum.—A number of lines in the red and orange yellow
and one in the blue. No bands in the green. Red 686.3 uy, 674.7 up,
662.8 wu, 649.9 wu; orange-yellow 646.5 wu, 635.1 uu, 606.0 uu; blue
460.7 uy.

108 REACTIONS OF THE METALS

BARIUM, Ba. At. Wt. 137.37
Sp. Gr. about 4.0. M. Pt. 850° (?) C.

Occurrence.—Like strontium, barium is almost always found
associated with calcium, but only in small amounts. The most im-
portant barium minerals are: Witherite, BaCO3, orthorhombic,
ismorphous with aragonite; barite, or heavy spar, BaSO4, ortho-
rhombic, isomorphous with anhydrite; and the hydrous barium
aluminium silicate, harmotome, BaAleH2Sis0i5+4H20. Harmotome ~
crystallizes in the monoclinic system, and belongs to the class of zeolites.

REACTIONS IN THE WET WAY

1. Ammonia and Ammonium Carbonate: same as with calcium
and strontium.

2. Ammonium Oxalate: same as with calcium and strontium,
except that the barium oxalate formed is more soluble in water (1 gm.
dissolves in 2.6 liters of cold water), and is readily dissolved by hot
dilute acetic acid.

3. Phosphates of the Alkalies: same as with calcium.

4. Chromates of the Alkalies produce in neutral solutions of barium
salts a yellow precipitate of barium chromate (thus differing from cal-
cium and strontium),

Bat to CrO4” = BaCrOx,.

The table on p. 10 shows that HCrO,- is ionized only to between 0.1
and 0.2 per cent in 0.1 N solution. In the presence of a stronger acid, its
ionization becomes much Iéss, and the equilibrium

Ht+cCrO,- —@ HCrO; and 2HCrO,- — H.O+Cr.0;=

progresses farther in the direction left to right than in pure water. The equilib-
rium, however, also depends upon the concentration of CrO,- in the solution.
In the case of strontium chromate, enough CrO, ions are present in the satu-
rated aqueous solution so that the formation of the HCrO, takes place to a
considerable extent when a little acetic acid is added to the solution, and it is
easy in this way to prevent the precipitation of as much as 0.5 gm. of strontium
ions; barium chromate, on the other hand, is so much more insoluble that it
takes considerable acetic acid to have an appreciable effect upon its solubility.
Thus a small quantity of barium can be separated from quite a large quantity
of strontium by adding CrO¢ and a suitable quantity of acetic acid.

If a more highly ionized acid is present, such as hydrochloric acid, the barium
will not be precipitated as chromate, but by adding sodium acetate (cf. p. 46)
the concentration of the hydrogen ions can be reduced sufficiently to precipi-
tate even a small quantity of barium as chromate.

BARIUM 109

5. Dilute Sulfuric Acid produces, in even the most dilute solutions,
a precipitate of barium sulfate:

Bat *++S04- — BaSO,.

According to the table on'p. 21, 1 liter of pure water dissolves about
2.5 mgs. of barium sulfate. In a slight excess of sulfuric acid it is
much less soluble by virtue of the common ion effect (p. 45). As with
strontium and barium sulfates, the presence of an excess of hydrogen

- ions has a solvent action due to the formation of acid sulfate. This

effect is appreciable with acids such as hydrochloric or nitric acid, but
the solubility of barium sulfate is so slight that it requires treatment
with hot, concentrated sulfuric acid in order to get any considerable
quantity of barium sulfate into solution, and dilution with water causes
reprecipitation of barium sulfate:

BaSO, +H2S804 2 Ba(HS04)s.

Barium sulfate is partially converted, as a result of the mass action
effect, into more soluble barium carbonate by boiling with a concen-
trated solution of sodium carbonate:

BaSO.+ NaeCO3 —@ BaCO3 +Na2SQO4.

To make this decomposition quantitative, the barium sulfate
must be boiled with the sodium carbonate solution, filtered, treated
with a new portion of sodium carbonate solution, and the process re-
peated until the filtrate no longer gives a test for sulfate. The more
concentrated the sodium carbonate solution is, the more complete
will be the decomposition. The highest degree of concentration will be
reached by fusion of the barium sulfate with anhydrous sodium carbonate.

Consequently, to obtain a solution of barium ions from insoluble barium
sulfate, it is best to proceed as follows: Mix the solid with four to six times as
much calcined sodium carbonate and fuse the mixture in a platinum crucible.
Cool, boil the residue with a little water until thoroughly disintegrated, and
filter. Wash the residue with hot, normal sodium carbonate solution, until the
filtrate gives no test for sulfate ions, and then with a little water. Dissolve
the residue of barium carbonate in dilute hydrochloric, nitric or acetic acid.

If the product of the fusion were treated with considerable water, or if the
residue were washed at once with considerable hot water, the dissolved sodium
sulfate would react with the insoluble barium carbonate to form less soluble
barium sulfate. This is prevented, in accordance with the mass-action prin-
ciple; by keeping the concentration of the sodium carbonate sufficiently large.

6. Fluosilicic Acid produces a white, crystalline precipitate of
barium fluosilicate:

H2SiF.+BaCle =2HCl+ BaSiFs.

110 REACTIONS OF THE METALS

In order to effect complete precipitation, the solution must stand some —
time. Barium fluosilicate is difficultly soluble in water and dilute
acids, and insoluble in alcohol.

7. Absolute Alcohol dissolves neither the nitrate nor the chloride;
neither of these salts is deliquescent: 100 cc. 66 per cent alcohol (by
volume) dissolve 3.3 gms. of BaCle-2H20. |

8. Concentrated Hydrochloric Acid and Nitric Acid will precipitate
from fairly-concentrated barium solutions the chloride and nitrate
respectively. |

REACTIONS IN THE DRY WAY

Heated with sodium carbonate on charcoal before the blowpipe, the
barium compounds, unlike those of calcium and strontium, do not give
a brightly luminous mass, because the barium carbonate formed is not
- decomposed at this temperature into the infusible oxide and carbon
dioxide, but merely sinters together somewhat. Volatile barium salts —
color the non-luminous gas flame yellowish-green. ,\ The sulfate is
only slightly volatile in the hottest flame, and in the ordinary gas flame
it shows scarcely any coloration. In order to obtain the coloration,
it is best to change the sulfate into chloride, by reducing a small
particle on a platinum wire in the upper reducing flame to sulfide and
adding a little hydrochloric acid by means of a capillary tube. The
wire on then being brought into the flame will give the characteristic
flame coloration.

Flame Spectrum.—A number of deep green lines, weaker lines in
the orange-yellow part of the spectrum, and one blue line.

Orange-yellow (654.0 uy, 629.8 pu), (624.0 uu, 617.9 pu, 610.9 wp,
603.2 wu). Yellow 528.5, yellow triple line (576.9 uy, 572 wy, 564.8 up).
Green 553.5 wy, 534.7 wy, 524.3 wy, 513.7 wy, 500.0 wu. Blue, 484.7 pp.

Separation of Calcium, Strontium, and Barium

In the course of a systematic analysis these three metals are always
obtained in the form of their insoluble carbonates, either by precipita-
tion with ammonium carbonate from fairly dilute solution in the presence.
of ammonium chloride (cf. p. 103) or by fusion of the sulfates with
sodium carbonate.

A number of excellent methods have been proposed for the analysis
of the alkaline earth group. The characteristic reactions of barium,
strontium and calcium are so similar that difficulties are likely to arise
in any scheme whenever a small quantity of one of these elements is
present together with a large quantity of another; thus a precipitate

STE a ee

Ne ee, le ee be
Pai ee oe

TAS

BARIUM 111

caused by the presence of much strontium may be confused with one.
produced by a little barium. Two methods of analyzing the group will
be described.

TasBLeE III.—Anatysis or Groupe IV. Mernop A.

Solution may contain: Bat+, Srt+, Ca++, Mg++, K+, Nat, NH«t. Add
NH,OH and (NH;,)2CO;. Examine filtrate for Mg, K, Na according to Table I, p. 97, or
Table II, p. 99. (1). Dissolve carbonates of Ba, Sr and Ca, in 2-normal HNOs3.
Evaporate to dryness. Dissolve a part of the residue in a little water and add CaSO,
solution (2); (a) no precipitate is formed. Ca is present (3); (b) a precipitate forms
slowly, Sr is present and possibly Ca (4); (c) a precipitate is formed at once, Ba is
present and possibly Sr and Ca (5). Treat the remainder of the dry residue with
C,H;OH (6).

Residue: Sr(NOsz)2, Ba(NOs)2, Heat with solid | Solution: Ca(NOs3)2. Evap-
s NH.,Cl (8). Treat cold residue with C2H;OH. orate to dryness and test
in the flame. Brick-red
coloration shows Ca. (7)

Residue: BaClh. Test in | Solution: SrCl,. Hvaporate
the flame. Dissolve in| to dryness and test in
water, heat to boiling and| flame. Carmine-red
add HC.H;0.and K2CrO,,| coloration shows Sr. (9)
yellow precipitate of ‘

BaCrO, shows Ba. (9)

PROCEDURE ’

1. Concentrate the filtrate from Group III to small volume in a porcelain
dish, add 6N HCl to acid reaction and filter if the solution is not perfectly
clear. If sulfur runs through the filter paper, make a pulp by shaking small
pieces of filter paper in a bottle with water, add some of this pulp to the turbid
solution, stir or shake, and filter. Wash the residue till the volume of the
filtrate is about 50 cc., heat this solution to boiling, add NH,OH and
(NH,).CO;. Filter and examine the filtrate for Mg*t, Kt and Nat. The
precipitate may contain BaCO;, SrCO; and CaCQs.

2. Dissolve the carbonates in dilute nitric acid and evaporate the solution

- in a small porcelain crucible just to dryness, by heating with a small flame kept

in constant motion and heating very carefully until all the moisture and nitric
acid have been expelled. Take care not to overheat the residue, as this will
cause the conversion of the nitrates to oxides. Dissolve a small portion of the
residue in as little water as possible and add 5 cc. of calcium sulfate solution.

3. If no precipitate is formed, only calcium (and magnesium) can be present.

4. If a precipitate forms only after standing some time, barium is absent,
strontium is present and calcium may be present.

5. If a precipitate forms immediately, barium is present and possibly the other
members of the group as well. This preliminary test with calcium sulfate
solution is very useful. It must be remembered in making this test that 5 ce.
of saturated calcium sulfate solution contain only 3 mg. of calcium ions so that
a large precipitate cannot be obtained with this reagent even when much barium
is present; 10 mgs. of barium will give the maximum possible precipitate with
5 ec. of calcium sulfate reagent.

112 REACTIONS OF THE METALS

6. Treat the.remaining residue, which must be perfectly dry, with a little
absolute alcohol, stir with a glass rod and pour the alcohol through a filter which
has been wet with absolute alcohol, catching the filtrate in a small porcelain
crucible. If any calcium was present in the original residue of nitrates, some
of it will have been dissolved by the alcohol.

7. Evaporate the alcoholic solution to dryness, wipe out the crucible with
a piece of filter paper containing no calcium, fold the paper and fasten it to a
platinum wire. Burn the paper, moisten the ash with hydrochloric acid and
hold it in the non-luminous flame. The presence of calcium is shown by a
brick-red flame coloration.

8. If calcium is found to be present, repeat the treatment of the original
nitrate residue with absolute alcohol several times to dissolve out all the cal-
cium nitrate. Then mix the remaining residue with an excess of ammonium —
chloride and heat until no more ammonium fumes are evolved. By this
operation the nitrates are converted to chlorides. The nitrogen, which has a
negative polarity of three in ammonium chloride (having four negative valences
and one positive) is oxidized by the quinquevalent, positively-charged nitrogen
of the nitrate and the product of this oxidation and reduction is nitrogen. At
the same time the negatively-charged chlorine is oxidized by the nitrogen of
the nitrate, forming free chlorine and nitric oxide, the latter escaping partly
as such and partly oxidizing the nitrogen of the ammonium. The oxidation
of the ammonium chloride requires four positive charges of electricity; the
nitrogen of the strontium nitrate loses five positive charges if reduced to nitro-
gen and three if reduced to NO. Assuming a complete reduction to nitrogen,
the equation is

28r(NO;).+6NH,Cl =28rCl,+5N, 7 +Cl Tt +12H,0;
assuming the nitrate reduced only to nitric oxide, NO.
Sr(NOs)2-+2N HCl =SrCl,+4H20+2NO 1 +N. 7.

It is also easy to change the nitrate to chloride by evaporation with con-
centrated hydrochloride (cf. p. 57).

9. Treat the residue of chlorides with a little 66 per cent alcohol (by vol-
ume) exactly as described above and test the alcohol solution for strontium by
the flame test; a carmine-red coloration shows the presence of strontium. Wash —
any residual chloride with 80 per cent alcohol to dissolve out all the strontium

chloride and test the residue likewise in the flame for barium; a yellowish-green _

coloration shows barium. If there is sufficient residue, dissolve it in a little water
and a little acetic acid, heat to boiling and test for barium with potassium
chromate solution; a yellow precipitate at this stage of the analysis is es
proof of the peice of barium.

The above cpthad of analysis is based on the assumption that the
alkaline-earth carbonates were formed under conditions such that only
barium, strontium and calcium are present in the carbonate residue or
precipitate. If the ammonium carbonate is added under the conditions
described on p. 100, the magnesium will be precipitated as carbonate,
together with calcium, strontium and barium. The following method

BARIUM 113

of analyzing this group has been tested thoroughly by A. A. Noyes*
and his students and permits the identification of as little as 1 mg. of
any constituent in the presence of 500 mgs. of any other member of the
group.

Taste ITV.—ANALysis oF Group IV. Mernuop B.

Solution may contain: Bat+, Srt+, Mg++, K+, Nat, NHit+. Concentrate to
10 cc.; add 15 cc. (NH1)2COs reagent, or more if necessary, and an equal volume of
C.H,OH. Stir, let stand 30 minutes and filter. Test filtrate for Na+ and K* accord-
ing to Table I, p. 97. Dissolve the precipitate, which may contain BaCO;, SrCOs,
CaCO; and MgCO3-(NH3)2CO3,-4H20, in 6-normal HC2H;0. add NH,C2H;O0. and
K2CrO,. (1) .

Precipitate:| Filtrate: Sr++, Cat+, Mg++. Add NH,OH. Dilute to 65 cc.
BaCrO,. Dissolve} and add 50 cc. C:H;OH. Shake the solution with filter paper
in HCl. LEvap-| pulpand filter. (8). . :
orate to dryness.

- Test residue in Li
flame, treat with Precipitate:

Filtrate: Cat +, Mg++. Dilute to 150 ce.

4
? ,
y
:

3 CC. of 6-normal SrCrOx. Pour a
Minriwre os
(NH,)2COz; and
K2CrO, through
the filter and reject
this filtrate. Wash
with cold water
and dissolve the
residue of SrCO3
inatlitttle
HC.H;0.. Test
with CaSQx,. Slight

HC2H;02, 20 cc.
of 3-normal
N H. 4CoH AOD and
15 cc. of water.
Heat to boiling

and test’ with

K2Cr2O,4_ solution.
Yellow precip-
itate is BaCrQy,.

— (2)

precipitate, shows

Sris present. (4)

heat to boiling andadd (NH,4)2C20,. (5)

Precipitate:

CaC.0,4. Dissolve
in 5 ce. of 6-
normal HySO. and
add 10 cc. of
C.H;OH. White
precipitate
is CaSO... (5)

Filtrate: Mgtt.

Add NH,OH and
N. aoH P O14. Dis-
solve in 2-normal
HSO,, add 20 ce.
C.H;OH and. filter
af necessary. Add
NH,OH. Precipi-
tateis MgN HyPQO,.
(6)

lk et i ee a i ae
= Yo

PROCEDURE

1. Dissolve the fairly dry carbonates in hot, 6 N acetic acid and evaporate
the resulting solution just to dryness, taking care not to overheat the residue.
Moisten the residue with 3 cc. of 6 N acetic acid and dissolve it in 20 ce.
of water. Add 20 cc. of 3 N ammonium acetate solution, heat to boiling
and precipitate the barium as chromate by adding hot 3 N potassium
chromate solution, drop by drop, until no further precipitation takes place
and the solution is decidedly yellow in color. Boil gently for about two minutes
longer and filter through paper capable of holding very fine crystals.

2. If the-filtrate is not decidedly orange in color, showing the presence of an
excess of dichromate ions, add a little more chromate to see if the precipitation
of the barium is complete. Wash the precipitate thoroughly with cold water,
rejecting all but the first of the washings. To confirm the barium test, dissolve

“A Course of Instruction in the Qualitative Analysis of Inorganic Substances.

114 REACTIONS OF THE METALS

the precipitated barium chromate in a little hydrochloric acid, evaporate the
solution to dryness and try the flame test (ef. p. 110). If the flame test is incon-
clusive, it may be that the precipitate contained strontium chromate, due to
the presence of a large quantity of strontium. In such cases repeat the precipi-
tation with potassium chromate, using the same quantities of acetic acid and
ammonium acetate as before. By evaporation to dryness, the chromate is
_ almost entirely changed to chloride and chromic salt and in the second pre-
cipitation with chromate in acetic acid solution, no strontium chromate will be
deposited, because the quantity of strontium now present cannot be large enough
to come down with the barium a second time.

3. To test for strontium, make the filtrate from the first precipitation with
potassium chromate distinctly ammoniacal, dilute to 65 cc. and add 50 ce. of
alcohol. If necessary add a little more potassium chromate and make sure
that the solution remains distinctly ammoniacal. Shake the solution with
some filter paper pulp and then filter. Allow the filter to drain thoroughly, but
do not wash the precipitate. Place the filtrate aside for the calcium and mag-
nesium tests.

4, Pour through the filter, which presumably contains strontium chromate,
a hot mixture of 10 cc. ammonium carbonate reagent and 5 cc. of potassium
chromate reagent. This mixture serves to convert strontium and calcium
chromate to the less soluble carbonates. If barium chromate is present it is
unaffected by the treatment. Wash the residue with cold water until the wash-
ings are colorless, which causes the removal of any remaining magnesium.
Dissolve the carbonate on the filter in a little normal acetic acid and test the
solution for strontium with calcium sulfate (ef. p. 111).

5. To test for calcium, dilute the filtrate from the strontium chromate pré-
cipitate to about 150 cc., heat to boiling and add ammonium oxalate solution,
in small portions, until no further precipitation takes place. Filter the solution
while still hot. The calcium is precipitated as oxalate and the magnesium
remains in solution. To confirm the calcium test, dissolve the precipitated
calcium oxalate, ora part of it, in 5 cc. of 6 N sulfuric acid and add 10 ee. of
alcohol. A white precipitate of calcium sulfate shows calcium to be present.
If the analysis is properly conducted, very little strontium, if any, will be pre-
cipitated with the calcium as oxalate and practically none of it will dissolve in
the sulfuric acid. The calcium sulfate test is shown with 0.2 mg. of calcium
ions if the solution is allowed to stand ten minutes. Not more than 0.3 mg.
strontium will be present if the above directions are followed properly. Mag-
nesium does not interfere with the test.

6. To test for magnesium, make the filtrate from the calcium oxalate precipi-

*- tate strongly ammoniacal and add a little sodium phosphate reagent (cf. p. 95).

To confirm the magnesium, dissolve the precipitate in 5 cc. of 2 N sulfuric
acid, add 20 cc. of alcohol and stir well for two or three minutes. Filter off
any calcium sulfate that may form, and repeat the precipitation of the i
- slum as phosphate (cf. p. 98).

Traces of alkalies and alkaline earths are recognized best by means
of the spectroscope. At this place, therefore, a brief description of
spectroscopic analysis will be given.

BARIUM 115

Spectroscopic Analysis (Bunsen and Kirchhoff, 1865)

_ If a ray of white light is passed through a glass prism, not only
is the direction of the ray changed, but the white light is decomposed
into colors; it suffers dispersion. It
will be found that the red rays are
deflected least, while the violet rays
are deflected most. The picture ob-
tained—the spectrum—if projected on
a screen (Fig. 5), does not show the
colors sharply separated, but merging
into one another. Such a spectrum
“is called a continuous, or uninter-

rupted, spectrum. LHvery glowing solid Rea
or liquid body emits white light; the Yellow
spectra obtained in all such cases will Violet

be continuous ones. Glowing vapors
and gases behave quite differently.
They do not emit white light, but Fia. 5.
light composed of rays of definite
wave lengths, which are characteristic for each gas and for each
vapor. The light emitted from glowing vapors or gases, when
decomposed by the prism, yields on the screen a discontinuous or

interrupted spectrum. If the light is passed through a fine slit before
reaching the prism, the spectrum will be found to consist of a greater
or less number of colored lines which always appear in the same place
with any given substance, provided the prism or its position is not
changed. In order to determine the exact position of these lines, we

116 REACTIONS OF THE METALS

make use of the spectroscope of Bunsen and Kirchhoff (Fig. 6). Fig. 7
shows a cross-section of the apparatus.
s :

Fria. 7.

The substance to be examined is placed in the loop of a platinum
wire and introduced at A into the non-luminous gas flame, by means
of which it is volatilized. The rays of light pass through the slit
into the tube Sp, reach the prism, by which the rays are refracted

BARIUM | 117

into the telescope C (the collimator) and are observed at D. Upon
a glass plate at the end of the tube Sk is a transparent scale, which is
illuminated by a small flame at B. This tube is so tnelined toward
the face aa of the prism that the rays of light from this tube are totally
reflected into the tube C, and reach the eye of the observer; thus the
rays from the substance appear at a certain position on the scale.
As, however, the position of the lines depends upon the dispersive
power of the prism, and upon its angle of refraction, it is clear that the
position of the lines will be somewhat different in different spectroscopes.
_ As every ray has a definite wave length, it is better to give the wave
__ lengths of the rays which appear, rather than their position on the scale.
Wave lengths are expressed in millionths of a millimeter, called
5

micromillimeters, and designated as up.
The following values for the wave-lengths of the various lines in
the spectrum are taken from the most accurate measurements:

Ree PUMPING § 5 sia bs os be Rew en Rbs =795.0 up

SRP MMO Gr a bes eh eens seo Sw ae ales Rb, =781.1 up
j potassium. double line..-....2......... K = 166.9 mp
SL OO. Oe
: SEMA T So re. eee a poe eal s Li, =670.8 up
| RCA NO oe hi elon oad eae es Cd, =643.9 up
Orange-yellow lithium line................ Lig +=610.3 pp

Yellow sodium (middle of the double line)...Na =589.3 up

RaPOHEY CAICWINL, DIG. f. )s os 6k bate ade Ss as Cag =554.4 pu

MRSMR TEED TING 9. ihe 5. 0 yes ex wsatig Sac ¢ bes 83020 un

MAMIDC TO Ce eee a eg Cd, =508.6 pu

Roribe COMM HOF os ce ek oie eco heen sha Cdg =480.0 uu
DPIC Boe he obs Ob sd thle Mate Cd, =467.8 up

SPOR GRNAIEL PIMOS ogg enka) db eh sone eras eds Sr; ~=460.7 wu

RAPOUMISR UTNE igs So ain ele! Su Sap pane ied Csg =459.3 uy

OE oo SUSE aS ee ial Ree Cs, .=455:3. ap

Blue-violet indium line...............:... Ing.) = 451; Tae

WAGIEL PMDICUUIN TING. Go) eek es ek dood Rbg =421.5 up

WER RCRRIS SURI te PGs oaks he 60s Rb, =420.2 up

BEA TEI Bi GS cons eae Re headed sos Ing =410.1 pp

WMIDGAMIIIN TO Ko Sol oes slere ceic es Kg, =404.4 up

mp (i ranennorer Gln)... 6355200 asco wcns ss = 396.8 up

Let us assume that the foregoing lines were observed at the following
positions on the scale: K, at 17, Kg at 154, Li, at 32, Na at 50, Tl at 68,
Sr; at 101, In, at 111, and Ingat 149. |

* The cadmium lines can be seen distinctly only in spark spectra.

118

‘

REACTIONS OF THE METALS

If now, upon a rectangular system of coordinates, we plot the wave
lengths as ordinates and the corresponding scale divisions as abscissas,

and connect the points of intersection,

ad ae we shall get a curve* which expresses
"75" os the relation between the wave lengths

and the corresponding divisions on the
we scale (Fig. 9). We can, without serious
125-4 error, regard the portions of the curve

between two points of intersection as
wach proportional to the corresponding scale -
675- divisions.

Lia The use of the curve may be illus-
rd trated by a single example:
625- Suppose we observe at the division

60 of the scale a linex. What is the
gi corresponding wave length? The line
BT5- ge lies between the sodium and thallium
oe lines, which are of known wave lengths.

Tla
525-
500-
Hp

fi
450 Ing
ag me Kp
400 } | tT t I T T T T - —

20 40 50 60 70 80 90 100 110 120 130 140 150 160

Fia. 9,

The sodium line lies at division 50 and corresponds to 589.3 pu
The thallium line

68 535.0 pu

Consequently 18 scale divisions= 54.3 up

and lscale division=3.017 py

The line x is at 60=50+10 divisions, and these correspond to
Division 50=589.30 yu
10 divisions= 30.17 py

Division 60= 559.13 wu

* The curve is ordinarily drawn only for wave lengths lying between 800-400 pp
because only these light rays are visible to the eye. The rays of shorter wave
length than 400 pp are called ultra-violet rays and those longer than 800 py are
The former can be detected by their action upon photo-

known as ultra-red rays.

graphic plates, and the latter by their thermic effect or by their action upon
specially prepared photographic plates.

;
kL.

BARIUM 119
The scale divisions increase in their value as the wave lengths
diminish, so that we subtract the 30.17 from 589.3.

The spectra thus obtained are very beautiful, but usually of short duration.
More permanent spectra may be obtained with the aid of the Beckmann burner, *
Fig. 8, p. 116. From one-half to 1 cc. of the salt solution to be tested is placed in
the glass vessel G, the gas is lighted, and a good blast of air introduced at a. By
means of the air current, a little of the solution is carried mechanically over into
the burner in the form of spray, and thus the salt reaches the non-luminous flame.
In this way it is possible to get a fairly permanent spectrum by the use of only a
few milligrams of substance. There is then plenty of time to make the measure-
ments without having to stop and replenish the sample,

Measuring the Lines and Bands of the Spectrum

The correctness of a curve of wave-lengths, prepared as just
described, depends upon the accuracy with which the observed lines
or bands are measured. ‘To insure accuracy, all the better forms of
apparatus are provided with cross-hairs in the ocular, and the cross must
each time be made to cover a definite part of the line to be measured.
The choice of such position depends upon the construction of the slit
at the end of the spectroscope. If the apparatus is provided with an
unsymmetrical slit, i.e., one of which only one side is movable while the
other remains in a fixed position, then the reading should be made with
the cross-hairs coinciding with the edge of the band on the immovable
side of the slit. Such a position is shown in Fig. 10, in which it is
assumed that the right-hand side of the slit is immovable.

In the case of a symmetrical slit, i.e., one in which both sides of
the slit open and close symmetrically, the cross-hairs must be made to
meet in the center of the line to be measured (Fig. 11).

The measurements can be made more readily and more accurately
by the aid of Hilger’s Wave Length Spectroscope (Fig. 12). In this
apparatus the telescope and collimator are both fixed in position, but
the prism can be rotated by means of a cylinder upon which is engraved
a very exact wave-length curve (Fig. 13). The pointer of the cylinder
gives the desired wave length with an accuracy of +0.2 uy.

For the rapid adjustment of the apparatus, the ocular shown in
Fig. 14 is convenient. In this ocular there is a very fine, polished steel
point beneath the cross-hairs; it is lighted from the outside by means of
a small mirror so that it is seen very distinctly. This point is made to
coincide with the desired edge, or the middle of the line, and then only
the final adjustment is made with the cross-hairs.

* Cf. Z. phys. Chem., 11, 472.

120 ) ‘REACTIONS OF THE METALS

It is not possible to volatilize all substances in the gas flame. The.
non-volatile substances will not give any flame spectra, but by means of
the electric spark they will be volatilized enough to give spark spectra.
As the determination of their spark spectra furnishes the only method
by which the purity of many substances may be tested, the methods for —
the production of spark spectra will be outlined briefly.

The apparatus invented by Bunsen consists of two platinum wires
which are each attached at one end to conical pieces of charcoal. The
latter are soaked with a solution of the substance to be tested. The

Movable Movable

Y YW
J SAIN

Fic. 10.—Position with un- Fig. 11.—Position with sym-
symmetrical slit. metrical slit.

two carbon points are now placed opposite to one another, quite near
together, and the other ends of the platinum wires are connected with
an induction coil, which causes sparks to pass between the two carbon
points, volatilizing some of the salt. If now the sparks are viewed
through a spectroscope, a large number of lines will be seen, of which
only a part are produced by the substance itself. Some of the lines are’
caused by the carbon points and some by the air. In order to determine
which lines are caused by the substance that is being tested, the experi-
ment is first performed without using any of the substance, and the
spectrum thus obtained is either drawn or photographed. The experi-
ment is then performed with some of the salt impregnated in the carbon
points and the new lines in the spectrum will be caused by the substance.
Spark spectra can be more conveniently produced by means of the

BARIUM 121

so-called “ fulgurator ”’ of Delachanel and Mermet, as shown in Fig. 15.
‘The salt solution is in a test-tube, so that the slit of the spectroscope

cannot be contaminated with the spattered particles of the salt.

Massegeatl!
j
|

Lh
iS,

| LAAT
LT

ro)
fl 63
eee,
3 ee
=
385

i CT

FOTO TLL
il

Fria. 14.—Ocular with steel point

Fic. 13.—Cylinder (enlarged).
and mirror.

The small apparatus of H. Dennis, Fig. 16, is also very convenient
for producing spark spectra. The platinum wire z is fused in a glass
tube and ends in a point of Ceylon graphite, which extends up out of the
arm e of the apparatus. The upper pole is not shown in the illustration.

122 REACTIONS OF THE METALS

To fill the tube m with the salt solution of the substance to be tested,
sis taken away, and the apparatus is inclined to the left. The solution
is poured in at the upper end of m, when s is again introduced and shoved
down until it reaches nearly to the bottom. The apparatus is now
placed in a vertical position, whereby the liquid in e rises until it is
level with the lower end of s. On raising s the liquid will rise to the
upper edge of e. By means of electric sparks from the carbon point,
enough of the liquid is evaporated to give the desired spectrum. Care
should be taken to prevent any spattering
of the substance into the slit of the spec-
troscope.

A far better fulgurator has been devised
by E. H. Riesenfeld and G. Pfiitzer.* As
anode the solution of the salt itself is used,
as cathode a thin iridium wire which almost
touches the surface of the liquid. If an
electric are is produced between such elec-
oes trodes, nearly pure spectra are obtained
: without lines from the atmosphere. In this
way even the magnesium lines can be
identified. By flame spectra with the
Bunsen burner only 0.2 mg. Ca, 0.6 mg. Sr
and 14mgs. Ba can be detected with cer-
tainty, but with this fulgurator 0.002 mg.
Ca, 0.03 mg. Sr and 0.006 mg. Ba can be
detected when present in 1 cc. of water.

To examine gas spectra, small Geissler
tubes are used which contain the gas to be

Fra. 15. Fia. 16. detected in a dilute condition. Besides

flame spectra and spark spectra, there re-
main absorption spectra to be mentioned. If white light is passed
through a colored solution or gas, certain rays are absorbed by the
liquid, so that if the light is now examined in the spectroscope, —
these rays will be found to be lacking. We see a bright spectrum
broken by black bands (absorption bands) which are character-
istic for different substances. Thus solutions of permanganic acid,
neodymium, praseodymium, erbium, and many other substances give
characteristic absorption spectra. These absorption spectra are often
obtained from colorless solutions.

It is to be noted that an absorption gaseunten is often of quite
different appearance according to whether the solution is dilute or

* Ber., 1918, 3140.

BARIUM 123

concentrated. In measuring absorption bands the solution should
be diluted with solvent until the band appears in the form of as fine a
line as possible at the intersection of the cross-hairs, and further dilu-
tion should cause the disappearance of the band.

The same end is easily attained if, as R. Philip has suggested, the
concentrated solution is placed in a test-tube and covered with more

: a aa BERSE
PSSST SSS

s TT HHH
N SS

BSS SSS SS y
ESS SSSR MSS

Fic. 17.

of the solvent. The lowest layer then gives the spectrum of the con-
centrated solution and it becomes more and more dilute in the upper
portions.

_ In order to obtain any of the above spectra sharply defined, it is
necessary to have the spectroscope properly adjusted, i.e., the rays must
come parallel from the collimator tube into the telescope. The telescope

is removed and adjusted for parallel rays by focusing it upon some
distant object. It is then replaced, the prism removed, and the colli-
mator tube brought exactly opposite the telescope, so that the slit of
the former can be observed by the latter. The collimator tube must be
- lengthened or shortened until the picture of the slit is sharply defined.
The prism is now replaced and the scale tube adjusted until the scale
also can be seen clearly defined. The instrument is now ready for use.

The direct-vision pocket spectroscope, with an arrangement of
prisms as shown in Fig. 17, is very convenient for ordinary use

124 REACTIONS OF THE METALS

If it is desired to photograph a spectrum, the telescope tube is
replaced by a camera as shown in Fig. 18. Such an apparatus is
called the spectrograph. With glass prism and glass objective lens,
the spectrum can be photographed to about 350 yu. To make the
ultraviolet part of the spectrum visible, the objective lens and the prism
must be prepared of quartz and fluorspar, whereby it is possible to
photograph light rays of about 200 yu. For observing rays of still
shorter wave length, fluorspar alone is requisite, as it absorbs the ultra
violet rays to a less extent than quartz does.

To photograph the ultra red rays, specially prepared plates, as dis-
covered and prepared by W. Abney, are necessary.*

*For methods of quantitative spectrum analysis consult W. Hempel and R. L.
von Klemperer, Z. angew. Chem., 1910, 1756.

“aa

GROUP III. AMMONIUM SULFIDE GROUP

ALUMINIUM, TITANIUM, CHROMIUM, IRON, URANIUM, ZING,
MANGANESE, NICKEL, COBALT

ALUMINIUM, Al. At. Wt. 27.1
Sp. Gr. 2.56-2.67. M. Pt. =658.7° C.

Occurrence.—Aluminium occurs very extensively in nature, prin-
cipally in the form of silicates, of which the feldspars and micas with
their decomposition products are important examples:

Orthoclase (feldspar), KsAls(Sig0g)3; muscovite (mica),
KHoeAl3(Si04)3; kaolin (decomposed feldspar), H4aAleSizO9. Impure
kaolin is called clay.

Among the most important minerals which contain aluminium
may be mentioned cryolite, NagAlF¢; spinel, MgAleO4, or magnesium
aluminate, which crystallizes in the regular system and is isomorph-
ous with magnetite, Fe-Fe2O4, and chromite, Fe-Cr204; alunite,
KAI3(SO4)2(OH)6; aluminium hydroxide as hydrargillite, Al(OH)s,
monoclinic; bauxite, H4AleO5, and diaspore, HAIOg, orthorhombic.

Corundum, AlgOs, is next to the diamond the hardest mineral.
Pure, transparent corundum often has a. beautiful color and is classed
with the precious stones. Thus the ruby, sapphire, oriental emerald,
oriental topaz and amethyst are nearly pure corundum colored by a
little foreign oxide. Emery is an intimate mixture of corundum,
magnetite and iron sulfide. True topaz is Ale(F,OH)2S8i04, turquoise
is Alo(OH)3PO4-H20, and garnet is chiefly CagAle(Si04)s. .

Metallic aluminium is now made on a large scale by the electrolysis
of AlsOs dissolved in a bath of molten cryolite. It has a silver-white
color and is only slightly attacked by exposure to the atmosphere.
From the position of this element in the electromotive series, (p. 41)
the metal should be attacked readily by the atmosphere, and the reason
it is apparently unattacked has been proved due to its becoming
coated with a thin, firmly adherent, protective layer of oxide. Alumin-
ium is trivalent in all its compounds; it forms only one oxide, Al2Os,
which is amphoteric. The metal dissolves in acid to form an aluminium
salt and in caustic alkali to form a soluble aluminate. Aluminium
readily replaces the hydrogen of hydrochloric acid, but it dissolves

125

126 REACTIONS OF THE METALS

less readily in dilute sulfuric acid and becomes passive when treated
with nitric acid. One theory of the cause of passivity is the formation
of a closely adherent oxide film.

The action of aluminium upon dilute hydrochloric acid and upon
aqueous solutions of caustic alkali is expressed by the following ionic
equations:

2A1+6H* — 2Al***+3H:2 T,
2A14+20H-+2H20 > 2Al02-+3He Tt.

Aluminium salts are as a rule colorless, and those which are soluble
in water show an acid reaction in aqueous solution, on account of their
being hydrolyzed to a considerable extent. This explains the fact
that on evaporating a solution of aluminium chloride in water we do
not obtain aluminium chloride, but the insoluble oxide, or hydroxide:

AICl3-+3HOH @ Al(OH)3+3HCI.

The property which aluminium possesses of forming alums is very
characteristic. The alums are double salts of aluminium sulfate with
the sulfates of potassium, cesium, rubidium, or ammonium, of the
general formula RAI(SO4)2:12H2O, in which R represents one of the
univalent metals just mentioned. Similar salts in which the aluminium
is replaced by trivalent iron or chromium are also called alums. The
alums crystallize in the regular system, usually in octahedrons. The
common potassium alum is much less soluble in cold than in hot water.
Thus 100 cc. of water at 15° C. dissolve 10.7 gms. of alum and 283 gms.
at 100° C.

The sulfide of aluminium can be prepared only in the dry way.
It is a pale-yellow substance, which is hydrolytically decomposed even
by cold water into hydrogen sulfide and aluminium hydroxide;

AloS3+6HOH =3HeS 7 +2Al1(OH)s.
Toward strong: acids aluminium hydroxide plays the part of an
alkali, while toward strong bases, on the other hand, it acts as an acid.
REACTIONS IN THE WET WAY

1. Ammonia produces a gelatinous precipitate of aluminium hydrox-
ide, which is somewhat soluble in water, but insoluble in the presence of
ammonium salts:

AICl3+3NHs0OH = Al(OH)3+3NH4Cl.

ALUMINIUM 127

The property which the aluminium hydroxide shows of partly
dissolving in water, is common to all colloidal substances (cf. p. 59).
When dissolved they are sometimes said to exist in the hydrosole con-
dition and when precipitated as hydrogele. :

The hydrosole form of aluminium hydroxide can be converted into
hydrogele by the addition of salts,* preferably ammonium salts. [Jf,
therefore, we desire to precipitate aluminium from a solution by means of
ammonia, we take care that ammonium chloride is present. }

The freshly precipitated aluminium hydroxide is readily soluble in
dilute acids; but after standing some time in a salt solution, or after
long boiling, it becomes more difficultly soluble, so that it is necessary
to digest it with acid for a long time in order to bring it completely into
- solution.

The solubility product of aluminium hydroxide is so small that it is
precipitated by ammonia even in the presence of ammonium salt.
Aluminium does not show a tendency to form soluble complex cations
with ammonia. |

2. Potassium or Sodium Hydroxide produces the same _ precip-
itate as ammonia, which is, however, in this case soluble in excess of
the reagent, forming an alkali aluminate:

Alt++430H- > Al(OH)s,
Al(OH)3+OH~ — AlO2~+2H20.

If dilute acid is added to a solution of an aluminate, there is formed
at first a precipitate of aluminium hydroxide, which dissolves on further
addition of acid:

AlO2- +H*+H20 — Al(OH)s,

Al(OH)3+3H+ > Al**+*+3H20.

The aluminates are also decomposed by boiling with an ammonium
salt,

AlOo-+NH,4t+H20 > Al(OH)3+NHs3 tT ‘

Aluminium hydroxide is soluble in neutral tartrates of the alka-
lies, so that in the presence of tartaric acid there will be no precipi-
tation on the addition of ammonia. Consequently the aluminium

* This principle is illustrated by the technical process of “ salting out ’’ colloidal
dyes from their solutions.:

128 REACTIONS OF THE METALS

cannot be present in the solution in the form of the simple alumin-
ium cation, but as a complex negative ion:

Al(OH)3+CsH406 — C4Heo(AlIOH)Og +2H20.

Many other organic hydroxy-acids and other hydroxy-compounds,
such as malic and citric acids, sugars and starches, have the same
effect of preventing the precipitation of aluminium hydroxide by am-
monia.

3. Ammonium Sulfide causes precipitation of the hydroxide.
When formed in the dry way, aluminium sulfide is comp'etely hydrolized
by water:

Al.S3-+6HOH =2Al(OH)3+3H2S 1.

The aqueous solution of ammonium sulfide is in equilibrium with
NH4OH and HeS formed by hydrolysis. There are, therefore, enough
OH ions in a solution of ammonium sulfide to cause the precipitation
of Al(OH)s.

4. Alkali Carbonates precipitate aluminium hydroxide (hydrolysis):

2AICls f 3NaeCO3 =6NaCl +Alo (COs3)s3,
Aleo(CO3s)3-+6HOH = 3H2CO3+2Al(OH)s3.
—
3(H20+COsz T )

5. Barium Carbonate, suspended in water and added to the solu-
tion of an aluminium salt, also precipitates the hydroxide:

2AICl3 +3BaCO3+6HOH =3BaCle+3H2CO3+2Al(OH)3.

—— .

3(H20+CO0z2 1)

6. Alkali Acetates produce no precipitation in cold neutral solu-
tions, but, on boiling the solution, a very voluminous precipitate of
basic aluminium acetate is formed:

AICl3 +3NaC2H302 =3NaCl+ Al(C2H302)3 (in the cold),

Soluble
SO me
_ Al(C2H302)3+2HOH @ mice +2HC2H30¢2 (on boiling).
CoH302

If the solution is allowed to cool, the basic aluminium acetate
redissolves. The reaction is, therefore, a reversible one, and goes
more completely from left to right according as we increase the amount
of water and raise the temperature (cf. pp. 50 and 51).

ALUMINIUM 129

7. Alkali Phosphates, e.g., NasHPOu,, give a gelatinous precipitate
of aluminium phosphate:

2HPO4« +Al**+ — AIPO,+HePOz,

or, on addition of ammonia,

HPO.-+NH3+Al*+t+ — AIPO,+NH,".

Aluminium phosphate is soluble in mineral acids, insoluble in
acetic acid (differing from Ca, Sr, Ba, Mg), but readily soluble in
sodium or potassium hydroxide solutions:

AIPO.,+40H — AlOo + PO4"+2H20.

On boiling this alkaline solution (obtained in the last reaction)
with ammonium chloride, a precipitate will be formed, consisting
of a mixture of aluminium phosphate and aluminium hydroxide;
while barium chloride, on the contrary, will precipitate barium phos-
phate from such a solution and leave the aluminate dissolved.

8. Sodium Thiosulfate, NazS2O0s,. completely precipitates the
aluminium as hydroxide on boiling:

2Al***+38203"+3HOH =38S+3S802 T +2A'(OH)s.

The sodium thiosulfate, being a salt of a weak acid, serves to neu-
tralize the H* ions formed by the hydrolysis of the aluminium salt, and
the free H2S203 is so unstable that it breaks down into S and SOz;
equilibrium is disturbed by the precipitation of the Al(OH)3 and §
and by the escape of SOz gas.

9. Morin in alcoholic solution shows a green fluorescence when
brought in contact with only a trace of neutral aluminium salt. (This
reaction is very sensitive.) Beryllium salts and salts of the rare earths
do not give the test.

10. Ether precipitates white, crystalline aluminium chloride,
AICl3-6H2O, from a concentrated solution which is saturated with
HCl gas. Aluminium may be separated from beryllium in this way.

DETECTION OF ALUMINIUM IN THE PRESENCE OF ORGANIC SUBSTANCES
WHICH PREVENT THE PRECIPITATION BY THE ABOVE REAGENTS

The presence of tartaric acid or other non-volatile, organic hydroxy-com-
pound prevents the precipitation with above reagents. To detect the presence
of aluminium in such cases, the organic substance must be first destroyed. This
is best accomplished as follows: Add some sodium carbonate and a little

130 REACTIONS OF THE METALS

potassium nitrate to the solution, evaporate to dryness in a platinum dish, and
ignite the residue, whereby the aluminium becomes aluminate and the organic
substance is destroyed with separation of carbon. Treat the residue with dilute
nitric acid and filter; the aluminium goes into solution as nitrate and can
be precipitated with any of the above reagents.

If sufficient nitrate was present, the carbon will be completely burnt to CO,,
and the residue then often contains undecomposed nitrate or nitrite. It would,
therefore, be unwise to treat the residue in the platinum dish with hydrochlorie
acid, as aqua regia would be formed and the platinum would be dissolved. Con-
sequently the residue is treated with nitric acid (or with hydrochloric acid in a
porcelain dish). :

When strongly ignited, aluminium hydroxide loses water and
forms the anhydride, Al2O3, which is scarcely soluble at all in hydro-
chloric and nitric acids. In hot, concentrated sulfuric acid, with
a little water, it will dissolve after long digestion. The ignited alumin-
ium oxide, as well as the natural corundum, is most readily brought
into solution by fusion with potassium pyrosulfate. The fusion is
accomplished in the following way: Take twelve times as much fused
potassium acid sulfate as there is oxide to get into solution and heat it
by itself in a large platinum crucible over a small flame. The acid
potassium sulfate melts readily, at about 300° C., gives off water
(causing frothing), and becomes changed into potassium pyrosulfate:

2KHSO4 => HeO0O+ KoSe07.

As soon as the frothing has ceased, the transformation into potas-
sium pyrosulfate is complete. Add the dry aluminium oxide to the
crucible and continue heating until the melt begins to solidify (showing
that a considerable amount of potassium sulfate, which melts much
more difficultly than the pyrosulfate, has been formed), then raise the
temperature and continue heating until the oxide has dissolved clear in
the melt. By heating the pyrosulfate, SO3 is given off, which at the
high temperature is very active:

K28207 = Ke804+80s3.

If the heating is too rapid, much of the SOsz is lost.
After the reaction is finished, the melt contains the aluminium
as aluminium sulfate in the presence of potassium sulfate,

3Ko8207+ AloO3 = Alo(SO4)3-+3K280a,

and both of these substances can be brought into solution by treating
them with water.

ALUMINIUM 131

The ignited aluminium oxide can also be brought into solution
by fusion with caustic alkalies:

AlzO3+2KOH = 2KA102+H20.

This last operation is usually carried on in a silver crucible, never in
platinum, as the latter would be strongly attacked.

Native AlzO3 (corundum, ruby, sapphire, emery) can be completely
brought into solution by fusion with potassium pyrosulfate.

REACTIONS IN THE DRY WAY

Aluminium compounds, on being heated with sodium carbonate
on charcoal before the blowpipe give a white, infusible, brightly glowing
oxide, which, when moistened with cobalt nitrate solution and again
heated, becomes a blue infusible mass (Thénard’s blue) :

AleO3 +CoO = CoAleQOx..

In carrying out this test it is extremely important not to use an
excess of cobalt nitrate, for this salt leaves black cobalt oxide behind:on
ignition and when an excess is present it entirely obscures the blue color
of cobaltous aluminate.

The test is usually applied to a precipitate of aluminium hydroxide.
The best way to carry out the test is as follows: Dissolve the precipitate
in 5 ec. of 6 N nitric acid and add to the solution half as many small
drops of 1 per cent cobalt nitrate solution as there are presumable
milligrams of aluminium present. Evaporate the solution nearly to
dryness, add a few drops of water and moisten a roll of filter paper
with the concentrated solution. Heat over the flame until the paper
is entirely consumed and then very strongly.

Thénard’s blue is infusible. If a fused mass is obtained in this test,
the presence of aluminium is doubtful, because all fused glasses, such as
borax beads, sodium phosphate beads, etc., are colored blue by cobalt
oxide.

Aluminium salts are not volatile, and do not color the flame. By
ignition in the air, all aluminium salts, with the exception of the phos-
phate and silicate, are decomposed, leaving behind the oxide:

PAICI3 +30 = AloO3+3Cle T ,
2Al(NO3)3=Als03+3N20s 1 ,
Alo(SO4)3 = Als03+3803 T .

132 REACTIONS OF THE METALS

CHROMIUM, Cr. At. Wt. 52.0
Sp. Gr.=6.81. M. Pt.=1510° C.

Occurrence-—Chromium occurs in nature as chromite, FeCreOs,
isomorphous with spinel (see aluminium); as monoclinic crocoite,
PbCrO4; and as laxmannite, a double compound of lead-copper phos-
phate and basic lead chromate, (Pb,Cu)3(PO4)2-PbsO0(CrO4)o. Fur-
thermore, it is found in small quantities in many silicates, such as
muscovites, biotites, augites, etc., and consequently in their weathering
products, as in many kaolins, bauxite, etc.

Metallic chromium is a white, crystalline powder. With oxygen
it forms the following oxides: chromous oxide, CrO; chromic oxide,
Cr2O3; chromic acid anhydride, CrO3; and chromium peroxides
corresponding to CrO4, CreOg9, CreO11 and CreOj3 (ef. p. 139).

The oxides CrO and Cr2O3 are basic anhydrides, and, on being
dissolved in acids, yield the corresponding salts, the chromous and
the chromic compounds. Chromium trioxide is the anhydride of the
hypothetical chromic acid, HeCrO4 and forms chromates with bases.
The chromium peroxides have never been obtained pure; salts of per-
chromic acid have been isolated and analyzed. (Cf, p, 139.)

I. Chromous Compounds

Chromous oxide is known only in the form of its hydroxide, Cr(OH)s,
which, on being dried, loses hydrogen and water, leaving behind chromic
oxide:

2Cr(OH)2 = He +H2O + CreOz.

Like chromous hydroxide, all chromous compounds are extremely
unstable, being changed readily by contact with the air into chromic
compounds. Only the halogen compounds, the phosphate, carbonate, and
acetate, are known in the dry state; the sulfate only in solution. ‘The
acetate, Cr(C2H302)2+H20, is a reddish-brown, crystalline substance,
insoluble in water, but readily soluble in hydrochloric acid. This
solution, as well as that of all chromous salts, absorbs oxygen with
avidity, and is consequently used in gas analysis for the determination
of oxygen in gas mixtures. Solutions of chromous compounds) are
obtained by the reduction of chromic compounds with nascent hydrogen
(zine and acid), out of contact with the air.

On account of the instability of these compounds the ariabplical
chemist will rarely meet them and further description is unnecessary.

CHROMIUM 133

II. Chromic Compounds

All chromic compounds contain chromium as a trivalent element;
they are colored either green or violet, and are soluble in water as
a rule. The oxide, hydroxide, phosphate, anhydrous chloride, and
the sulfate, after being strongly heated in a stream of carbon dioxide

gas, are insoluble in water. Violet chromium chloride obtained in

the dry way is insoluble in acids; it dissolves readily in water containing
a trace of chromous chloride, or in the presence of stannous chloride
(tin and hydrochloric acid). By dissolving the grayish-green chromic
hydroxide in acids, green solutions are always obtained, which on
long standing become greenish violet or violet, but on boiling become
green again. Chromic sulfate forms with sulfates of potassium,
ammonium, cesium, or rubidium, the so-called chrome-alums, which
crystallize in the regular system. ‘These alums, like all other chromic
salts, react acid in aqueous solution (hydrolysis).

Chromic sulfide, CreS3, can be obtained only in the dry way.
On being treated with water it is decomposed quantitatively into
hydroxide and hydrogen sulfide.

REACTIONS IN THE WET WAY

1. Ammonia produces a erayish-green, gelatinous precipitate of
chromic hydroxide:
Crt +++30H- — Cr(OH)s.

Chromic hydroxide is somewhat soluble in excess of ammonia, form-
ing a violet or pink solution, particularly soluble when the ammonia
is added to a violet solution of a chromic salt, in the presence of am-
monium salts. This is caused by the formation of complex chromic-
ammonia cations, which, however, may be decomposed by boiling
the solution until the excess of ammonia has been driven off, when the
chromium is quantitatively precipitated as hydroxide. In order, then,
to precipitate the chromium from a solution as hydroxide, it is neces-
sary to precipitate at a boiling temperature, and to use as little
ammonia as possible. |

By ignition of chromic hydroxide, green chromic oxide is obtained, which
after strong ignition is insoluble in acids. In order to bring it into solution, it
must be fused with potassium pyrosulfate (cf. aluminium, p. 130); or with
an oxidizing flux such as sodium carbonate and potassium nitrate in a platinum
crucible or sodium peroxide in a nickel crucible, whereby sodium chromate is
formed:

| Cr.0;+2Na.CO;+30 =2Na,Cr0O,+2CO, T “3

134 REACTIONS OF THE METALS

If the product of this last fusion is dissolved in water, acidified with hydro-
chloric acid, and boiled with alcohol, a green solution of chromic chloride will
be obtained (p. 188), from which the chromium can be precipitated as
hydroxide with ammonia. On fusing with sodium carbonate and potassium
nitrate in a platinum crucible, the crucible will always be slightly attacked, so
that a small amount of platinum will go into solution with the fused mass; it
can be removed, after the treatment with hydrochloric acid, by passing hydrogen
sulfide into the boiling solution, and filtermg off the precipitated platinum
sulfide.

2. Potassium and Sodium Hydroxides cause the same precipitation
as ammonia; but the precipitate is readily soluble in excess of the
reagent, forming a green chromite:

Crt +t+4+30H— —> Cr(OH)s3; Cr(OH)3+OH 2 CrO.-+2H20.

Chromic hydroxide behaves here as a weak acid. The. reaction
is reversible, the presence of considerable water causing the reaction
to go from right to left, particularly at the boiling temperature. By
boiling the dilute solution, complete hydrolysis takes place; the chro-
mium is almost quantitatively precipitated as hydroxide (differing from
aluminium).

Chromic hydroxide often causes the precipitation of other bases
as insoluble chromites, particularly zinc and magnesium.

3. Sodium Peroxide and caustic alkali added to the solution of a
chromic salt converts the trivalent chromium into the chromate ion |
in which the chromium has a positive valence of six:

2Cr* ++43Na002+40H- — 2CrO4-+6Nat+2H.0.

4, Alkali Carbonates, Barium Carbonate, Ammonium Sulfide,
and Alkali Thiosulfates precipitate chromic hydroxide, as with
aluminium.

5. Alkali Phosphates give a greenish, amorphous precipitate of
chromic phosphate:

2QHPO4-+ Crt t+* — CrPO4+HePO,".

Chromic phosphate is readily soluble in mineral acids and in cold
acetic acid. On boiling.the acetic acid solution, chromic phosphate
separates out again.

6. Alkali Acetates produce in solutions of chromic salts no pre-
cipitation, even when the solutions are boiled. If, however, consider-
able amounts of aluminium and ferric salts are present at the same
time, the chromium will be precipitated almost quantitatively with the
iron and aluminium as basic acetate. In case, however, chromium
predominates, only a part of the metals will be precipitated as basic

CHROMIUM | 135

salts; the filtrate will contain iron and aluminium with chromium.
In the presence of. chromium, the basic acetate separation is always uncer-
tain.

7. Hydrogen Sulfide produces no precipitate in acid solutions
of chromic salts.

III. Chromates

Chromium trioxide, CrOzs, forms red orthorhombic needles, which
dissolve readily in water to an orange-red solution. If this solution

is neutralized with potassium hydroxide, it becomes yellow, and on

evaporation yellow KeCrO,, the potassium salt of chromic acid, H2CrOu,
is obtained. If the yellow solution of potassium chromate is acidified,
and then allowed to crystallize, orange-red prisms of triclinic potassium
dichromate crystals, Ko2Cr207, are formed.

The aqueous solution of potassium chromate, K2CrO4, contains
colorless potassium ions and yellow CrO4~ anions, while the aqueous
solution of potassium dichromate, KeCreO7, contains, in the presence
of the colorless potassium ions, the orange-red colored Cr207~ anions:

KoCre07 = 2Kt + CreO7- -

We are able, therefore, to determine from the color of a chromate
solution the nature of the chromate ion which is present.

The free acids, H»CrO, and H.Cr,O;, cannot be isolated, but only the corre-
sponding anhydride, CrO;, which is very soluble in water. When chromium
trioxide dissolves in water, the following reaction takes place:

CrO; +4- H.O @ H.CrO,.

The first hydrogen of this acid belongs to the class of very strong acids (cf
p. 10) and consequently the greater part of the H.CrO, undergoes ionization
as fast asitisformed: H,CrO,— H*+HCrO,. Thesecond hydrogen, on the

other hand, corresponds to that of a very weak acid and we may say that in

the presence of hydrogen ions HCrO,~ is scarcely ionized at all. The HCr0,-,
however, enters into the following equilibrium: 2HCr0O,- 2 H.O+Cr O;.
The presence of hydrogen ions will evidently favor the formation of the dichro-
mate ion, while dilution will favor the reverse reaction, and the presence of
OH ions will cause the dichromate ions to disappear and yellow CrQ,- ions
will take their place.

Remark.—Although we may judge as to the color of the ions from the color
of the solution, and often predict what the color of the solid salt will be, yet, on
the other hand, we cannot tell what the color of the solution will be from that
of the salt itself. Yellow lead iodide dissolves in water to a colorless solution,
and the yellow and red iodides of mercury, although only silghtly soluble, also
do not yield colored solutions.

If the solution of a salt is colored, the salt itself will be colored; but the reverse
ts not always true.

136 _ REACTIONS OF THE METALS

All chromates are insoluble in water, except those of the alkalies,
calcium, strontium, and magnesium. All chromates dissolve in nitric
acid, except fused lead chromate, which dissolves with difficulty.

Formation of Chromates ‘

All chromium compounds may be readily oxidized to chromates.
According to whether the compound is soluble in water or not, different
methods are used to effect the oxidation.

The student should not attempt to memorize a large number of chemical
equations, but he should strive to become able to express his chemical knowledge
in the form of equations or’ equilibria expressions. To balance equations
representing the oxidation of chromium, it should be remembered that the
chromium is changed from a positive valence of three (in the form of chromic
cations or chromite anions) to a positive valence of six (in the form of chromic
acid, chromate or dichromate ions). When the oxidation takes place by means
of halogen, a halide is formed and the halogen is changed from the neutral
condition to the form of a negative ion with unit charge. When hypochlorite
is the oxidizing agent, the unit positive charge on the chlorine atom is lost and
a unit negative charge takes its place, which corresponds to the loss of two
positive charges. Similarly, lead peroxide and hydrogen peroxide have an oxidiz-
ing power corresponding to the loss of two unit charges of positive electricity.

The valence of the chromium in the anions CrO,-, CrO;- and Cr,O,- is
found, in accordance with the rule given on p. 25, by subtracting the charge
of the ion from the product obtained by multiplying the number of oxygen
atoms present in the ion by its valence of two. In equilibrium expressions it
is important to make sure that the algebraic sum of the positive and negative
charges on one side is exactly the same as the algebraic sum of the charges on
the other. With a little practice it is very easy to determine whether hydro-
gen ions, hydroxyl ions or water molecules are required to make the equation
balance.

The oxidation in alkaline solutions is effected:

(a) By the halogens. If sodium or potassium hydroxide is added
in excess to a solution of a chromic salt, and chlorine or bromine is
conducted into the solution, the oxidation will be complete in a few
minutes: the green chromite becomes yellow chromate:

2Cr02~+80H-+3Cle > 2CrO4"+6CI--+4H20.

Chromic compounds may be also oxidized by halogens in the
presence of sodium acetate, the reaction going extremely slowly in the
cold, but very quickly on warming:

2Cr* *++4+3Cle+8H20 & 2CrO4-+6Cl-+ 16H.

The sodium acetate greatly lowers the concentration of the hydrogen
ions and permits the reaction to proceed slowly in the direction left to
right (cf. p. 46).

CHROMIUM 137

(b) By hypochlorites (sodium hypochlorite, chloride of lime, ete.):
2CrO2-+30Cl +20H- — 2Cr04-+3Cl-+H20.

2CrOe +3PbO02+80H — 3PbO2 +2Cr04-+4H20.
(d) By hydrogen peroxide,
2CrO2 +3H202 +20H” — 2CrO4 +4H20,

the reaction taking place on warming.
(e) By freshly precipitated manganese dioxide. The oxidation
_ takes place on boiling the neutral or slightly acid solution:

2Cr* +*+4+3Mn02+40H- — 2Cr047+3Mn* t+2H20.

It is evident from this last equilibrium expression that the presence of.

hydroxyl ions should favor the oxidation and hydrogen ions should hinder it.
On the other hand, manganese dioxide is very insoluble in alkaline solutions.
It is necessary, therefore, to have the solution neutral or slightly acid in order to
obtain a sufficient concentration of quadrivalent manganese in solution.

In acid solutions the chromic cation is the most stable condition for chrom-
ium, but in alkaline solutions the chromate anion is the more stable condition.
In acid solutions, therefore, it is easy to reduce a chromate to chromic salt and
in alkaline solutions it is easy to oxidize a chromic salt to chromate.

Oxidation in acid solution may be effected by boiling with very
energetic oxidizing agents such as concentrated nitric acid and potassium
chlorate, sodium bismuthate (or bismuth tetroxide), or potassium per-
manganate:

2Cr* +++3NaBi0g+4H* —> Cre07-+3Nat+3Bit +++2H20.

In carrying out this reaction the chloride should not be used, or any other
salt of which the anion is capable of oxidation; the acid used should
be nitric or sulfuric acid.

In the case of an insoluble chromium compound, such as strongly
ignited chromic oxide, or the mineral chromite, the oxidation is effected
by means of fusion with sodium carbonate and an oxidizing agent such
as potassium nitrate, potassium chlorate or sodium peroxide (cf. pp.
133, 142). The alkali chromates thus obtained are of a deep-yellow
color, and are readily soluble in water.

a Reduction of Chromates

Chromic acid, chromates and dichromates are strong oxidizing
agents in acid solutions. Such reactions often take place even in very

138 REACTIONS OF THE METALS

dilute solution, and for this reason potassium dichromate is often used
in quantitative analysis, the quantity of reducing agent being de-
termined by the volume of potassium dichromate required to react
- with it. Ferrous ions, sulfurous acid, hydrogen sulfide and hydriodie
acid are oxidized at the ordinary temperature. Oxalic acid and alcohol
are oxidized slowly at the laboratory temperature and very quickly
on heating the solution; hydrochloric acid and hydrobromic acid only
when the solution is hot. The original orange solution is changed to
green, the color of chromic ions;

Cr207-+6Fe* ++14H* > 2Cr*++*++6Fet*+++7H20,
Cr207- +3803" +8H*t — 2Cr++*+3S804-+4H20,
Cr207-+3H2S+8H*t — 2Crt +*+3S8+7H20,
Cr207~+3H2C201+8H* => 2Cr+*+*++3C02 T +7H20,
Cr2O7-+6I- +14H* > 2Cr*+++3I1,+7H:20,
Cr207~+6HC1+8H*t — 2Cr+*++3Cle T +7H20.

As this last reaction takes place only on warming, it furnishes us
with a convenient method for preparing small quantities of chlorine
for analytical purposes, because the evolution of chlorine ceases as —
soon as the lamp is taken away. It is necessary, however, to em-
ploy an excess of hydrochloric acid, as otherwise no chlorine will be
evolved owing to the formation of potassium chlorochromate, KCrOs3Cl:

Cre07~+2HCl — 2CrOsCl + H20,
which is decomposed on adding more hydrochloric acid:

2CrO3Cl-+4HCI+8H*t > 2Crt+*+*++6H20+3Cl. T.

If alcohol and hydrochloric acid are allowed to act simultaneously
upon a chromate (the reaction takes place on gentle warming without
the evolution of chlorine), the alcohol is oxidized to aldehyde:

Cr207"+3C2Hs0H+8H* — 2Cr***+-7H20+ 3CHsCHO hy
Idehyde
This last reaction is often used for reducing a chromate, because the
aldehyde (recognizable by its peculiar empyreumatic odor) and the
excess of alcohol are easily removed by boiling the solution, and the
latter then contains simply the chromium and the metal of the chromate
as chlorides.

a ec «
: Se te
J = |

ee ee a a a

CHROMIUM 139

By boiling chromates with concentrated sulfuric acid, reduction
takes place with evolution of oxygen:

2KeCr207+8H 2804 = 2K2804+ 2Cre(SO4)3 +8H20+302 T.

The behavior of free chromic acid toward hydrogen peroxide is
characteristic. The chromic acid is converted into blue perchromic
acids which are soluble in ether: H7CrOi0, H3CrO7 or HgCrOsg.

If a cold, alkaline solution of a chromate is treated with neutral
hydrogen peroxide, the solution is colored red, owing to the formation
of an alkali salt of perchromic acid, H3CrOs:

2KeCrO4+7H202+2KOH =8H20+2K3CrOsg.

Little by little the red color disappears, with evolution of oxygen,
and the yellow color of the chromate returns:

4K3Cr0g+2H20 =4KOH+4KoCr04+702 fT.

Tf a cold, neutral solution of potassium dichromate is treated with
hydrogen peroxide, the solution is colored violet, due to the formation
of the potassium salt of a slightly different perchromic acid, H3CrO7:

KeCre07 + 5H2O0e2 = 3HeO -- 2K HeCrO7z.

In this case, also, the violet color gradually disappears with evolution
of oxygen and regeneration of the dichromate:

4K He2CrO7 = 2KeCreO7 +502 T a 4H>0.

If either the red or violet solution, obtained as above described,
is shaken with ether, the latter remains colorless.

The behavior of chromate solutions toward an excess of hydrogen
peroxide in the presence of dilute sulfuric acid is quite different.
There is then formed invariably the perchromic acid richest in oxygen,
H7CrQjo, and the solution is turned an intense blue. The blue color dis-
appears after a short time and the solution turns green, owing to the
conversion of all the chromium into the chromic condition:

2H7Cr019+3H2804 = Cro(SO4)3-+10H20+502 T .

The perchromic acid is very soluble in ether; if, therefore, the
aqueous solution is shaken with ether the latter becomes colored a
beautiful blue. The perchromic acid is more stable in ethereal than
in aqueous solution.

Since the formation of the intense-blue-colored perchromic acid
takes place so readily, it may be used as a basis for a sensitive test

140 REACTIONS OF THE METALS

for free chromic acid, which is made as follows: Add a few drops of
dilute sulfuric acid to one or two cubic centimeters of hydrogen per-
oxide and shake with 2 cc. of ether; then add a little of the chromate
solution and shake the mixture again. In the presence of 0.1 mg. of
chromic acid, the upper ether layer is colored intensely blue, and ‘the
reaction is noticeable with only 0.007 mg. of chromic acid (Lehner).

Most chromates are insoluble in water, and exhibit character-
istic colors; therefore it is easiest to test for chromium when it is present
as a chromate.

Reactions for the Precipitation of Chromic Acid

1. Sulfuric Acid.—Dilute sulfuric acid causes, at the most, a
change of color from yellow to orange, without any evolution of gas.

Concentrated sulfuric acid causes the cold solution to change
to orange color, and there is often a separation of red needles of CrO3;
the solution on being heated becomes green, the chromic acid being
reduced to chromic salt with evolution of oxygen:

4Cr03+6H2804 =6H20+302 7 + 2Cro(SOx)s.

. 2. Silver Nitrate produces in neutral chromate solutions a brownish-
red precipitate of silver chromate:

CrO4g” + 2Agt —s AgoCrO4 )

soluble in ammonia and mineral acids (hydrochloric acid changes
it into insoluble silver chloride and chromic acid), insoluble in acetic
acid. If to a moderately concentrated solution of potassium dichro-
mate, silver nitrate be added, a reddish-brown precipitate of silver
dichromate is formed: .

CroO7_+ 2Agt — AgeCreOz,

which, on being boiled with water, is changed into the less soluble
normal silver chromate:

2AgeCr207 +H:20 — 2AgoCrO, + HeCre2QO7.

The presence of sodium acetate causes this change to take place in
the cold (cf. pp. 46 and 135).

3. Lead Acetate produces in solutions of normal chromates and
dichromates a ye'low precipitate of lead chromate, which is soluble
in nitric acid but insoluble in acetic acid:

CrO4” of Pbt ty PbCr O4

CHROMIUM 141

and

Cr207~+2Pb(C2H302)2+H20 — 2HCeH302+ 2C2H3027 + 2PbCrOg.

‘If lead nitrate is used instead of lead acetate, the precipitation is not
complete unless sodium acetate is added.

4. Barium Chloride produces in solutions of normal chromates
a yellow precipitate of barium chromate:

CrO4-+Ba*+t — BaCrOu,

soluble in mineral acids, insoluble in acetic acid. From solutions
of dichromates the precipitation is complete only on addition of ‘an
alkali acetate (cf. p. 108).

5. Mercurous Nitrate produces in the cold a brown, amorphous
precipitate of mercurous chromate:

CrO4” + Hgo(NO3)2—-Hg2Cr04+2NO037,

which on being boiled becomes fiery-red and crystalline.

Behavior of Chromium Trioxide and Chromates on Ignition

Chromium trioxide is decomposed on ignition into chromic oxide
and oxygen, 4CrO3=2Cr203+302 T. The chromates of ammonium
and mercury behave quite similarly. Thus normal ammonium chromate
on ignition is changed to chromic oxide, ammonia, nitrogen, and water.
The reduction of the chromate is favored by the reducing action of
ammonia which is present in excess.

2(NH4)2CrO4=2HN3 7 +Ne2 1 +5H20 Tf +Cr20s.
Ammonium Dichromate evolves only water and nitrogen:
(NH4)2Cr207=4H20 T Nz T +CreOs.

This decomposition takes place violently with scintillation. The
chromic oxide which remains behind is very voluminous and reminds
one of tea-leaves; consequently it is sometimes called ‘‘ tea-leaved

chromic oxide.”’ :
Mercurous Chromate is decomposed on ignition into chromic oxide,
mercury vapors, and oxygen:

AH g2CrO4 =2Cr203-+8Hg 1 +5027.

The Dichromates of the Alkalies are changed on ignition into normal
chromates, chromic oxide, and oxygen:

4KoCre07 =4KoCr04+ 2Cre03+302 t P

142 REACTIONS OF THE METALS

REACTIONS OF CHROMIUM IN THE DRY WAY

All chromium compounds color the borax, or salt of phosphorus, —
bead an emerald green both in the oxidizing and reducing flames.
Heated with sodium carbonate on charcoal before the blowpipe, all
chromium compounds yield a green slag, which after long heating is
changed to green infusible chromic oxide. By fusing with sodium
carbonate and potassium nitrate in the loop of a platinum wire, all
chromium compounds yield a yellow melt of alkali chromate:

2Cre03+4Na2C03+302 = 4NaeCr04+4CO0e ; .

If the fused mass is dissolved in water and acidified with acetic acid,
the solution will give with silver nitrate a reddish-brown precipitate of
silver chromate. This reaction is very delicate and serves for the
detection of minute traces of chromium. Cloth which has been dyed
with a chromium mordant can be tested in this way; the ash from a
thread 5 cm, long is sufficient to give the test.

IRON, Fe. At. Wt. 55.9
Sp. Gr. 7.98. M. Pt. about 1600° C.

Occurrence.—Native iron is rarely found. It occurs in basaltic
rocks; also in meteorites, associated with nickel, cobalt, carbon,
sulfur, and phosphorus.

The most important iron ores are the oxides and the sulfides.
Of these may be mentioned:

Hematite, Fe2O3, isomorphous with corundum; magnetite, Fe30x,
isomorphous with spinel; géthite, FeHO2, isomorphous with diaspore
and manganite; limonite, Fe4H¢O9; (bog ore), Fe(OH)s, which is
used in the purification of illuminating gas; pyrite, FeSe, which crystal-
lizes in the isometric system; marcasite, FeSe, orthorhombic. Iron
disulfide is, therefore, dimorphous. Another important iron ore is
siderite, FeCO3, which is rhombohedral; vivianite, Fe3(PO4)2-8H20,
is monoclinic.

The metallic iron of commerce is never pure, but usually contains
more or less iron carbide, iron sulfide, iron phosphide, iron silicide,
corresponding manganese compounds and graphite, etc.

On dissolving commercial iron in acids (H2SO4,HCl), hydrogen,
contaminated with small amounts of hydrocarbons, hydrogen sulfide,
mercaptans, phosphuretted hydrogen, and silicon hydride is given off,
and these impurities give to the gas its unpleasant odor. There
remains almost always an undissolved residue consisting chiefly of
carbon.

|
4
7

TRON 143

Tron is bivalent, trivalent, and rarely hexavalent, forming the
following oxides:

Iron protoxide Iron sesquioxide Ferrous-ferric oxide eg
or ferrous oxide or ferric oxide or magnetite Iron trioxide
FeO Fe203 Fe304 FeOs.

Iron trioxide, FeO3, containing hexavalent iron, has never been
isolated. It plays the part of an acid anhydride in ferrates of the general
formula R2FeO4, which are decomposable by water.

By dissolving these oxides in acid the corresponding salts are
obtained; thus ferrous oxide gives with hydrochloric acid, ferrous

: chloride,

FeO0+2HCl=H20+ FeCl;

ferric oxide gives ferric chloride,
FeeO3 + 6HCl = 3HeO te 2FeCls ’

while ferrous-ferric oxide yields a mixture of ferrous and ferric chlorides:
Fe304+8HCl =4H20+2FeCl3+FeCle.

Iron, therefore, forms two-series of salts: first, the ferrous, derived
from ferrous oxide, containing bivalent iron; second, the ferric, derived
from ferric oxide, containing trivalent iron. These two series of salts
show a quite different behavior toward reagents.

The position of iron in the electromotive series shows that it is capable of
being oxidized by hydrogen ions and suggests the possibility of its being
oxidized to a measurable extent by the hydrogen ions of water. Careful.
experiments have shown that absolutely pure water containing no dissolved
oxygen, does not affect iron appreciably at ordinary laboratory temperatures,
though there is evidence that a trace of iron dissolves and a film of hydrogen
is formed on the metal which acts as a check upon further attack. The position
of ferrous iron in the voltage series shows that hydrogen ions cannot oxidize
iron appreciably to the ferric condition. The presence of dissolved oxygen,
and this is normally present in all water that is exposed to the atmosphere,
can accomplish this oxidation of the ferrous ions to the ferric condition and it
also aids in the oxidation of the iron from the metallic to the ferrous state. In
this way iron exposed to moisture and oxygen oxidizes or rusts. The rusting
process is favored by the contact of the metal with a more noble metal, such as
platinum, copper or nickel; an electric couple is formed and the iron becomes
the positive pole, so that the hydrogen set free by the action of iron upon water
is deposited upon the more noble metal. The presence of a less noble metal,
such as zine, tends to hinder the corrosion of iron: the zinc corrodes instead of
the iron.

The presence of an acid is, therefore, not absolutely necessary to start the
corrosion of iron. An increase in the concentration of hydrogen ions, however,
will greatly hasten the solution of the metal. When carbonic acid is present,
ferrous bicarbonate is first formed and, when the ferrous iron is oxidized to the

144 REACTIONS OF THE METALS
4

ferric condition, the carbonic acid is set free again because ferric carbonate does
not exist. The acid again acts upon the metal and the rate of its corrosion is —
greatly accelerated.

Certain substances tend to make iron passive, particularly strong nitric acid.
Passive iron does not dissolve in dilute nitric acid and does not corrode readily
(cf. aluminium, p. 126). On the other hand, certain substances can overcome
the passive condition and are said to activate the iron. Thus a solution of com-
mon salt is an activating agent.

Different varieties of iron and steel corrode with different degrees of readi-
ness. Cast iron is often protected by its casting skin. Impurities present in
steel often favor corrosion by causing electric couples to be established,

A. Ferrous Compounds >

Ferrous compounds, which may be prepared by dissolving metallic
iron, ferrous oxide, ferrous hydroxide, ferrous carbonate, or ferrous
sulfide, ete., in acids, are usually greenish in the crystallized state,
but in the anhydrous condition they are white, yellow or bluish; in
concentrated solution they are green; in dilute solutions almost color-
less. Ferrous compounds exhibit a strong tendency to change over into
ferri¢ salts; they are strong reducing agents.

‘REACTIONS IN THE WET WAY

1. Ammonia produces in neutral solutions an incomplete pre-
cipitation of white ferrous hydroxide:

FeCle-+2NH3+2H20 @ Fe(OH)2+2NH4Cl.

Ferrous salts in this respect are similar to those of magnesium
(cf. p. 94). In the presence of ammonium chloride the reaction takes
place in the direction from right to left; ammonia, therefore, causes
no precipitation with ferrous salts out of contact with the air, provided
sufficient ammonium chloride is present. On exposure to the air,
however, a turbidity is soon formed, green at first, then almost black,
and finally becoming brown. The small amount of ferrous hydroxide
contained in the solution is oxidized by the air, forming at first black
ferrous-ferric hydroxide and finally brown ferric hydroxide.

2. Potassium and Sodium Hydroxides produce, if air is excluded,
complete precipitation of white ferrous hydroxide,

Fe+20H~ — Fe(OH)s2,

which is quickly oxidized by the air into ferric hydroxide.

3. Hydrogen Sulfide produces no precipitation in acid solutions
of ferrous salts; in dilute neutral solutions a small amount of black
ferrous sulfide is precipitated; but if the solution contains consid-

IRON 145

erable alkali acetate, hydrogen sulfide precipitates more of the iron as
ferrous sulfide (but not all of it), in spite of the fact that ferrous
sulfide is readily soluble in acetic acid. This interesting fact is an
instructive illustration of the law of chemical mass action.

The table on p. 22 states that 3.41078 -g. of ferrous sulfide dissolves
in a liter of water. This small quantity exists in solution almost entirely as
Fet* and S~ ions. FeS —@ Fet*++S>. When acetic acid, which is a much
stronger acid than hydrogen sulfide (cf. p. 10), is added to the solution, equilib-
rium has to be established between its hydrogen ions and the dissolved sulfide
ions, 2H*+S~ — H.S, and, as a result of the formation of non-ionized hydrogen
sulfide, the solution no longer contains enough sulfur ions to reach the value of
the solubility product of FeS; to restore the equilibrium between FeS and its
ions, more of the solid must dissolve. If, moreover, the solution is boiled, the
hydrogen sulfide escapes as a gas as soon as it is formed. Consequently it is
impossible to arrive at a state of equilibrium until all of the ferrous sulfide has
dissolved. The solution is accomplished by means of hydrogen ions:

FeS+2Ht = Fet+t+H,8 [.

On the other hand, if the ionization of the acetic acid is repressed by adding
an alkali acetate to the solution (cf. p. 46), and the concentration of the hydro-
gen sulfide is made as large as possible by keeping the solution saturated with
the gas, the reaction will take place in the reverse direction and some of the
iron will be precipitated as ferrous sulfide.

4. Ammonium Sulfide precipitates iron completely as black fer-
rous sulfide:

FeClo+(NH4)2S = 2NH4Cl+FeS,

which is readily soluble in acids with evolution of hydrogen sulfide.
In moist air it turns slightly brown, a part of the sulfur separates out,
and a basic ferric sulfate is formed.

5. Alkali Carbonates precipitate the white carbonate,

FeCle +NaeCO3 =2NaCl +FeCOs,
which in contact with the air becomes green, then brown:
4FeCO; +6H20+02 =4CO>2 +4Fe(OH)s,

being converted into ferric hydroxide with loss of carbonic anhy-
dride. |

Ferrous carbonate, like calcium carbonate (cf. p. 103), is soluble in
carbonic acid, forming ferrous bicarbonate:

FeCO: + HeCO3 = FeHo(CQOs) 2;

a compound which is found in many natural waters, but which, like
the normal carbonate, is decomposed by atmospheric oxygen with
separation of ferric hydroxide:

4eHs (COsz) 2 +2H:2O +02 = 8CO2 +4Fe(OH)s3.

146 REACTIONS OF THE METALS

Consequently a mineral water which contains ferrous bicarbonate,
if allowed to stand in contact with the air, will become turbid, owing
to the deposition of ferric hydroxide. To prevent this, the bottle
must be filled with water and tightly corked, so that no trace of air
can get in. Ferric hydroxide is insoluble in carbonic acid.

6. Potassium Cyanide precipitates yellowish-brown ferrous cyanide,

Fet++2NC” — Fe(CN)o,

which is soluble in excess of the reagent, forming potassium ferro-
cyanide:
Fe(CN)2+4CN” — Fe(CN)6 ~.

The complex ferrocyanide anion is in equilibrium, to be sure, with
simple ferrous cations, but the quantity of the latter present in the
aqueous solution of a ferrocyanide is so small that none of the above
reactions characteristic of ferrous ions can be obtained with it. Many
other similar complex cyanide anions are known; thus, the cyanides
of silver, nickel, iron (ferrous and ferric), and cobalt all dissolve in potas-
sium cyanide, forming the following: complex ions: [Ag(CN)e]",
[Ni(CN)4]~, [Fe(CN).6] =, [Fe(CN)6]~~ [Co(CN)e]>. The acids are:
H[{Ag(CN)2], He[Ni(CN)4], Hs[Fe(CN)6], HalFe(CN)6], Hs[Co(CN)sl.
It is possible, as a matter of fact, to isolate the last three acids, though
the two former have never been prepared; they immediately break
down into metallic cyanide and hydrocyanic acid, just as carbonic acid
is decomposed into water and carbon dioxide.

With iron, therefore, there are two series of complex cyanogen
compounds, the ferrocyanides and the ferricyanides. The ferro-
cyanic derivatives contain the quadrivalent ferrocyanide anion and the
ferricyanides contain the trivalent ferricyanide anion.

Potassium ferrocyanide, K4{[Fe(CN).], is often called yellow prussiate
of potash, and potassium ferricyanide, Ks3[Fe(CN)g¢|, is called red prus-
siate of potash. The solubility of the alkali and alkaline-earth salts, and
the insolubility and color of the salts of the heavy metals (especially
with both ferric and ferrous iron), are very characteristic of ferro- and
ferricyanides.

7. Potassium Ferrocyanide, K,Fe(CN)», produces in solutions
of ferrous salts, with complete exclusion of air, a white precipitate of
potassium ferrous ferrocyanide or of ferrous ferrocyanide, depending
upon whether one or two molecules of ferrous salt react with one mole-
cule of potassium ferrocyanide:

K4[Fe(CN)¢]-+FeSO4 = KoS04+KoFe[Fe(CN).]
K,[Fe(CN)¢]+2FeSO4 = 2K2S04+Fe2[Fe(CN)¢]

IRON : 147

Although both of the above salts are white, a light-blue color is almost
always obtained, because the precipitate is immediately oxidized some-
what by the air, forming the ferric salt of hydroferrocyanic acid (Prus-
sian blue):

3Fe2[Fe(CN)¢6]-+-3H20+30 =Fe,[Fe(CN )elg3 +2Fe(OH)3.

8. Potassium Ferricyanide, K3[Fe(CN).«], added to solutions of
ferrous salts produces a dark blue precipitate (Turnbull’s blue) consist-
ing of ferrous ferricyanide mixed with potassium-ferric ferrocyanide:

2K3[Fe(CN).6]-+3FeCls =6KCl+Fes[Fe(CN).]2

and
Ks[Fe(CN).6]+FeCle = KFe|[Be (CN) 6]~~+2KCl.

in other words the ferricyanide acts both as a precipitant and as an oxidizing
agent * and a blue color results whenever iron is present in the cation in a state
of oxidation different from that of the iron present in the complex anion. The
ferricyanide ion is a strong oxidizing agent and in alkaline solution readily
oxidizes ferrous hydroxide to ferric hydroxide. Turnbull’s blue is not very
soluble in acid solutions, but is decomposed by treatment with caustic alkali,
all of the complex anion being in the form of ferrocyanide:

Fe;+ tiFe(CN)ele =4 SKOH = 2K,[Fe(CN)<| +2Fe(OH);+Fe(OH),,
Kt¥Fet + t+[Fe(CN).] ==+3KOH =K,[Fe(CN).] +Fe(OH);.

9. Potassium Thiocyanate gives no reaction with ferrous salts
(note difference from ferric salts).

As has been stated, ferrous salts are oxidized by the air to ferric salts; thus |
ferrous sulfate is gradually changed into brown, basic ferric sulfate,

2FeSO.+0 =Fe.0 -(SOx)s,

’ which is insoluble in water. Consequently it often happens that ferrous sul-
fate will not dissolve in water to a clear solution, but gives a brown, turbid
solution, becoming clear on the addition of acid, the basic ferric salt being
changed to a soluble neutral salt:

Fe,0 : (SOx)2 +H.SO, == Fe, (SOx) 3+ H.0.

Such a solution, which then contains ferric ions, reacts with potassium thio-
cyanate (cf. p. 150). To free the solution from ferric salt, it may be boiled
with metallic iron, with exclusion of air, whereby the ferric salt is changed into
ferrous salt:

Fe.(SO,);+Fe =3FeSO,.

By means of strong oxidizing agents, ferrous salts can be quickly and
completely changed into ferric salts, as was shown in the introduction (cf.
pp. 27-33).

* Cf. Erich Miuer, J. pr. Chem., 84 (1911), 353.

148 _ REACTIONS OF THE METALS

Detection of Ferrous Oxide in the Presence of Metallic Iron

Treat the mixture with a large excess of a neutral solution of mercuric
chloride and heat on the water-bath; the metallic iron goes into solution as
ferrous chloride:

2HgCl, 4. Fe= FeCl, + Hg,Cl.

Filter off the residue and test the filtrate with potassium ferricyanide; a
precipitate of Turnbull’s blue shows that metallic iron was originally present.

Wash the residue with cold water, until all of the ferrous chloride has been
dissolved, and then treat it with dilute hydrochloric acid. If the solution
now gives a precipitate of Turnbull’s blue with potassium ferricyanide, ferrous
oxide was present. ‘

If hydrogen is given off, some metallic iron is still present; the experiment
must be repeated and the mixture given a longer treatment with HgCl, solution.

B. Ferric Compounds

Ferric oxide, Fe2O3, is reddish brown, becomes grayish black on
strong ignition, but on being pulverized appears red again.

The ferric salts are usually yellow or brown, but ferric am-
monium alum is pale violet. Ferric salts are yellowish brown in aque-
ous solution, and the solution reacts acid (hydrolysis). Dilution and
warming favor the hydrolysis, so that all strongly diluted ferric salts
deposit basic salts on being boiled:

Fee(SO4)3+H20 @ Fe2(SO4)20+ H2S80..

With ferric salts of the weaker acids, often all of the iron is pre-
cipitated as a basic salt; thus the acetate, on being boiled in a dilute
solution, reacts as follows:

Fe(C2H302)3-+2H20 @ Fe(OH) 2(C2H302) +2HC2H302.

By the addition of acid all basic salts may be changed back into
neutral salts.

REACTIONS OF FERRIC SALTS IN THE WET WAY
- 1. Ammonia precipitates brown, gelatinous ferric hydroxide:
Fet*++30H- — Fe(OH)3.

The solubility product of ferric hydroxide is so small (ef. p. 22)
that it is precipitated completely even in the presence of ammonium
salts; it is readily soluble in acids. On ignition it loses water and is
changed to oxide, which is very difficultly soluble in dilute acids. It is
best brought into solution by long-continued heating below the boiling
point with concentrated hydrochloric acid.

{RON 149

2. Potassium and Sodium ee also precipitate ferric hy-
droxide.
3. Sodium Carbonate lice a brown precipitate of basic car-

_ bonate, which at the boiling temperature is completely decomposed

. >

_hydrolytically into hydroxide and carbon dioxide:

2FeCl3+3Na2CO3+3H20 =2Fe(OH)3+6NaCl+3C02 Tf.

4, Zinc Oxide and Mercuric Oxide also precipitate the iron as

hydroxide:

2FeCls3+3Zn0+3H20 =3ZnCle+ 2Fe(OH)s.

_ This reaction is frequently used in quantitative analysis.

5. Sodium Phosphate precipitates yellowish-white ferric phosphate:
FeCl3+2Na2HPO4 =3NaCl+ NaH2P01+FePOx,.

Ferric phosphate is insoluble in acetic acid, but readily soluble
in mineral acids. The precipitation of iron with sodium hydrogen
phosphate is consequently only complete when a large excess of the

‘precipitant is employed, or when sodium acetate.is added:

FeCls +NasHPO, +NaCe2H30e =3NaCl + HC2H302 +FePO..

In this last case all the iron and all the phosphoric acid are pre-
cipitated. The reaction is often used to precipitate phosphoric acid
quantitatively. An excess of the disodium phosphate will also cause —
complete precipitation of iron as phosphate, if the phosphate solution
is previously exactly neutralized with ammonia:

NazHPO4+NH40H = H20+NazNH4POu,
and
NaeNH4P04+ FeCls = 2NaCl+ NH4Cl+FePO..

If, however, an excess of sodium phosphate and ammonia is added
to the iron solution, the precipitation of iron is incomplete, because the
ferric phosphate dissolves in the excess of sodium phosphate, in the
presence of ammonia (or ammonium carbonate), with a brown color
and formation of a complex salt.

Ferric phosphate is transformed by ammonia into a brown basic
phosphate, and by potassium hydroxide almost completely into ferric
hydroxide and potassium phosphate; while by fusion with caustic
alkali or alkali carbonate it is completely decomposed.

If alkaline earth ions are present, an excess of ammonia completely
changes ferric phosphate to ferric hydroxide and alkaline carth phos-
phate is precipitated.

150 REACTIONS OF THE METALS

6. Alkali Acetates produce in cold, neutral solutions a dark-brown
coloration, and on boiling the dilute solution all of the iron separates
as basic acetate:

FeCl3+3NaC2H302=3NaCl+Fe(C2H302)3 (in the cold), -
Fe(C2H302)3-+2H20 = 2HC2H302+Fe(OH)2C2H30z (on boiling).

The presence of organic hydroxy-acids (tartaric, malic, citric, ete.)
and of polyatomic alcohols (glycerol, erythritol, mannitol, sugars, etc.)
prevent all of the above-mentioned reactions, because complex salts
are formed in which the iron is present in the form of a complex anion

(cf. aluminium, p. 128).
‘7. Potassium Thiocyanate, KCNS, produces in solutions of ferric

salts a blood-red coloration:
Fet+++3CNS~ = Fe(CNS)s.

This action is reversible; the red color of the slightly ionized ferric thio-
cyanate being most intense when an excess of ferric salt, or of potassium thio-
cyanate is present.

If the solution is shaken with ether, the Fe(CNS); goes into the ether. Ferric
thiocyanate combines readily with potassium thiocyanate, forming complex
potassium ferrithiocyanate:

Fe(CNS);+3KCNS =K,[Fe(CNS)].*

analogous to potassium ferricyanide, K,[Fe(CN)s].

The complex salt is insoluble in ether, the Fe(CNS); only being soluble
therein; so that the red color is due to the formation of the ferric thiocyanate
and not to the complex salt.

This reaction is extremely sensitive, but cannot always be relied on. The
test cannot be made in the presence of strong oxidizing agents such as nitric
acid, as a red color is produced by the oxidation of the thiocyanate. The
oxidized compound is not very stable, however, and its color is not very deep.
If the solution contains considerable alkali acetate, the coloration cannot be
recognized. The presence of organic hydroxy-compounds (tartaric acid, ete.)
prevents the reaction in neutral solutions, but not in acid solutions. In the
presence of mercuric chloride the red color disappears entirely; the mercuric
chloride reacts with the ferric thiocyanate, forming a colorless, soluble
mercuric double salt, which is ionized even less than ferric thiocyanate:

2Fe(CNS);+6HgCl, =2FeCl,;+3[Hg(CNS).-HgCl.].
8. Potassium Ferrocyanide, K4Fe(CN).6, produces in neutral or
acid solutions of ferric salts an intense blue precipitation of Prussian

blue: :
3[Fe (CN) 6] ee 4Fet tt, Fe, [Fe (CN) 6ls-

* K;[Fe(CNS)«]+4H.0. Cf. RosenueErm, Z. anorg. Chem., 27 (1901), 208.

IRON 151

Prussian blue, the ferric salt of ferrocyanic acid, is insoluble in water, but
soluble in oxalic acid and in an excess of potassium ferrocyanide; the solution
thus obtained is a deep blue and is used as blueing and as blue ink. The blue
solution obtained with a ferric salt and an excess of potassium ferrocyanide con-
tains colloidal KFe[Fe(CN).]-H.O, which can be salted out by the addition of a
considerable quantity of electrolyte such as alkali chloride, sulfate or nitrate.
Prussian .blue is also soluble in concentrated hydrochloric acid, but is pre-
cipitated again on dilution. As the ferric salt of ferrocyanic acid it behaves
like other ferric salts to the hydroxides of the alkalies; ferric hydroxide and the
alkali salt of hydroferrocyanic acid being formed:

Fe,[Fe(CN).];+120H- — 4Fe(OH);+3[Fe(CN),|=.

9. Potassium Ferricyanide, Ks|Fe(CN).»], produces no precipita-
tion in solutions of ferric salts, only a brown coloration (differing
from ferrous salts) :

[Fe(CN).6]=+Fet ** = 2Fe(CN)s3.

10. Ammonium Sulfide, added to a solution of a ferric salt, gives
a precipitate of ferric sulfide, Fe2Ss,

2Fet $y T439= <> Fe2Ssz,

which is soluble in cold, dilute hydrochloric acid, forming ferrous
chloride and sulfur:

Fe.S3+4H+t > 2Fet++2H.S fT +S.

The fact that Fe.S; is precipitated, and not FeS as commonly believed, was
proved by H. N. Stokes* who decomposed it out of contact with air by zinc-
ammonium oxide and obtained white ZnS and red Fe(OH);._ L. Gedelt has also
shown that hydrogen sulfide passed into a solution of ferric chloride made
alkaline with ammonia gives Fe.8,. If, however, the solution is acid, hydrogen
sulfide or ammonium sulfide reduces the iron before any precipitate is formed.

11. Hydrogen Sulfide in acid solutions reduces ferric salts to
ferrous salts, with separation of sulfur:

2Fe*++t++H.S — 2Fet*+2Ht-+S.

Besides hydrogen sulfide, many other substances (nascent hydrogen,
stannous chloride, sulfurous acid, hydriodic acid, etc.) will reduce ferric
salts, as was shown on pp. 35, 36.

12. Ether when shaken with a solution of ferric chloride in 6 N
hydrochloric acid dissolves most of the ferric chloride. By separating

* J. am. Chem. Soc., 29 (1907), 304.
t Ueber Schwefeleisen, Karlsruhe, (1905).

152 REACTIONS OF THE METALS

the ether with the aid of a separatory funnel, and repeating the opera-
tion, nearly all of the iron can be removed from the aqueous solution.
(CH, 7... 1%)

13. Cupferron, the ammonium salt of phenylnitrosohydroxylamine,
CesHsNO-NONHg, precipitates red (CeHsNO-NO)s3Fe, which is soluble
in ether, insoluble in acids, and converted into Fe(OH)3 by treatment
with ammonia.

14. Sodium Thiosulfate, NazSeOQ3, colors neutral ferric solu-
tions a violet red, but the color disappears quickly and the solution
then contains ferrous salt and sodium tetrathionate;

2Nao8003+-2FeCls = 2NaCl+2FeClo+NaeSi0c.

The composition of the violet-red substance which is first formed
is unknown; perhaps it is ferric thiosulfate.

As we have seen, there exist a number of iron compounds which contain
the metal as a complex ion, so that it cannot be detected by the ordinary
reagents. The complex hydroxy-organic compounds, as well as the ferro-
and ferricyanide compounds, belong to this class of compounds.

If it is a question of proving the presence of iron in such a compound, a
different method should be used in the case of an organic hydroxy-compound
from that in the case of a ferro- or ferricyanide.

If organic substances are present, the iron is precipitated as sulfide by means
of ammonium sulfide; or the organic matter is first removed by ignition,
whereby metallic iron, ‘oxides of iron and carbon are obtained.

In case we have a ferro- or ferricyanide, the iron cannot even be piecinieeae
by means of ammonium sulfide; the compound must be completely destroyed
before it will be possible to detect the peer? of iron by any of the ordinary
methods.

This may be accomplished (a) by ignition, (6) by fusion with potassium |
carbonate or sodium carbonate, or (c) by heating strongly with concentrated
sulfuric acid.

(a) Decomposition by Ignition.—The ferrocyanides are decomposed (with
evolution of nitrogen) into potassium cyanide and carbide of iron;

K,[Fe(CN). =4KCN+FeC.+N, tT .

The ferricyanides also leave behind iron carbide and potassium cyanide,
but evolve cyanogen as well as nitrogen:

2K;[Fe(CN).] =6KCN+2FeC,+(CN).+2N, fT.

Treat the residue from the ignition with water, whereby the potassium
cyanide goes into solution, leaving behind the iron carbide; filter and treat
- the residue with hydrochloric acid. The iron goes into solution as ferrous
chloride, hydrocarbons are given off, and there remains some carbon.

The above decomposition can be imagined to take place as follows:

IRON 153

By heating potassium ferrocyanide, it is decomposed first into potassium
cyanide and ferrous cyanide, while the latter on further heating is changed
to iron carbide and nitrogen:

(a) K.[Fe(CN)«] =4KCN +Fe(CN):;
(8) Fe(CN)2=FeC:+N, T.

Potassium ferricyanide is decomposed into potassium cyanide and the
- very unstable ferric cyanide, which splits off cyanogen and becomes ferrous
cyanide; the latter is decomposed, as before, into iron carbide and nitrogen:

(a) K;[Fe(CN).] =3KCN +Fe(CN);;
(8) Fe(CN);=Fe(CN).+CN T ;
(y) Fe(CN). =FeC, +N, t °

(b) Decomposition by Means of Fusion with Potassium Carbonate.—Mix the
substance with an equal amount of the carbonate and heat in a porcelain cru-
cible until a quiet fusion is obtained. By this means a mixture of potassium
cyanide and potassium cyanate (both soluble in water) is formed in the presence
of metallic iron:

K,[Fe(CN).]+K.CO; =5KCN+KCNO+CO, f +Fe.

Extract the melt with water, filter and dissolve the iron in hydrochloric
acid.

(c) Decomposition by Heating with Concentrated Sulfuric Acid.—By heating
with concentrated sulfuric acid all complex cyanogen compounds may be
decomposei. By this means the metal present is changed into sulfate, the
nitrogen of the cyanogen into ammonium sulfate, while the carbon of the
_ cyanogen escapes as carbon monoxide:

K,[Fe(CN).] +6H,S0,+6H,0 =2K,S0,+ FeSO,+3(NH,).S0,+6CO Tf ,
2K;[Fe(CN).] +12H,SO,+12H,0 =3K.80,+ Fe.(SO,); +6(NH,)2S0,+12CO Tf .

The treatment with concentrated sulfuric acid is best accomplished in a
procelain crucible placed in an inclined position over the flame, and the flame
directed against the upper part of the crucible. Continue heating until fumes
of sulfuric acid cease to come off. Treat the residue, which consists of an
alkali sulfate and anhydrous ferrous or ferric sulfate, with a little concen-—
trated sulfuric acid, heat gently, and add water little by little. In this way the
sulfate is readily brought into solution.

REACTIONS IN THE DRY WAY >

The borax (or sodium metaphosphate) bead, containing a small
amount of an iron salt, is yellow while hot and colorless when cold
after being heated in the oxidizing flame, and pale green after being
heated in the reducing flame. When strongly saturated, however, the
bead obtained with the oxidizing flame is brown while hot, yellow

» \

154 REACTIONS OF THE METALS

when cold; and after heating in the reducing flame it becomes
bottle-green.

Heated on charcoal with soda before the blowpipe, all iron com-:
pounds leave a gray particle of metallic iron, which is usually difficult
to see, but can be separated from the charcoal by means of the magnet.
The reduction on the charcoal stick, as described on p. 65, is a
much more delicate test.

URANIUM, Ur. At. Wt. 228.5
Sp. Gr. =18.33. M. Pt.=2500(°?)

Occurrence.—Uranium occurs in nature chiefly in the mineral
pitch-blende, U30g; but it is also found in a few rare minerals, uranite,
(UO2)2Cu(PO4)2+8H20; samarskite (a niobite of iron, yttrium,
cerium, and erbium with varying amounts of uranium); and liebigite,
U(COs)2-2CaCO3+10H20.

Klaproth showed, in 1789, that the mineral pitch-blende contained
a new metal, which he called uranium. By heating the oxide with
reducing agents he obtained a brown, almost copper-red, substance, ©
which he took to be the metal, and it indeed does behave like a metal,
dissolving in acids in contact with the air, forming yellowish-green salts.

It was not until 1842 that it was shown by Péligot that this reddish-
brown body was not the metal uranium, but its dioxide. The hexa-
valent metal itself was obtained by Péligot, as a gray powder, by reduc-
ing the tetrachloride with sodium.
| Out of contact with the air, uranium dioxide (uranyl) dissolves in

strong acids, forming wranous salts:

U02+4HCl=2H20+UCh, UO2+2H2804 = 2H20+ U(SO4)o.

The uranous salts are extremely unstable, and on being exposed
to the air change rapidly, forming wranyl salts which contain the
bivalent UOz group:

UClk,+0+H20 — UOeCle+2HCl,
U(SO4)2+0+He20 = U02804+ HeSOu.

Only the reactions of the uranyl salts will be described in this book.

Besides uranyl (or uranium dioxide) uranium forms a trioxide,
UOz, which can be regarded as uranyl oxide, UO2O. It dissolves
in acids, forming urany] salts:

U03+2HCl = UOe2Cle+H20.

- URANIUM 155

By igniting the oxides of uranium in air, dark-green urano-uranic
oxide, U3Og or (2U03-UQOz2), is obtained, which out of contact with
the air dissolves in strong acids, forming a mixture of uranous and
uranyl] salts:

(2U03-UO2) +4H2804 = 2U02804+ U (S04)2+4H20.
By dissolving in aqua regia, uranyl chloride is obtained:
3U30s+18HCl1+2HNO3=9U02Cl2+2NO 7 +10H20.

All uranyl compounds are colored yellow or yellowish green. Most
of them are soluble in water, but the oxides, the sulfide, phosphate,
and uranates are insoluble. In mineral acids all uranium compounds
are soluble, with the exception of the ferrocyanide.

REACTIONS OF URANYL COMPOUNDS IN THE WET WAY

1. Potassium Hydroxide precipitates yellow amorphous potassium
uranate. Uranyl hydroxide, UO2(OH)s, is first formed and changes
into uranic acid, H2U207, of which the alkali salts are insoluble:

2U02Clz+6KOH — K2U207+4KC1+3H20.
2. Ammonia precipitates yellow, amorphous ammonium uranate:
2U02 (NOs) 2 +6NH40OH os (NH4)2U207 +4NH4NOs +3H20.

The alkali uranates are soluble in alkali carbonates, particularly in ammo-
nium carbonate, with the formation of complex salts:

(NH,).U20;+6(NH,4)2CO;+3H20 — 2(N H,).{UO.(CO;)3]++6NH,OH.

Consequently, in the presence of sufficient alkali carbonate, ammonia fails to
precipitate uranium. ‘Tartaric and citric acids (and other organic substances)
also prevent the precipitation with ammonia and caustic alkalies, as with iron,
chromium, and aluminium. —

3. Sodium Carbonate produces in concentrated solutions an orange-
yellow precipitate of sodium uranyl carbonate:

UO2(NO3)2+3Na2CO3 = 2NaN03+Na4[U02(CO3)3].

Sodium uranyl carbonate is soluble in considerable water, so that
no precipitate is formed from dilute solutions. It is still more soluble
in alkali carbonate solution, particularly in a bicarbonate solution.
From such solutions sodium hydroxide precipitates sodium uranate,
but ammonia produces no precipitation.

156 REACTIONS OF THE METALS

4, Barium Carbonate precipitates in the cold all of the uranium, ral
probably as barium uranyl carbonate:

UOe2(NO3)2+8BaCO3 = Ba(NOz)2+Ba2{U02(COs3)3].
5. Ammonium Sulfide precipitates brown uranyl sulfide,
UO2(NOs)2+(NH4)2S = 2NH4NO3s+U0O2S,

soluble in dilute acids and in ammonium carbonate:

U028+3(NH4)2CO3 = (NH4)28+ (NH4)4[U02(COs)3].

Ammonium sulfide, therefore, produces no precipitate in solutions of
uranyl salts in the presence of ammonium carbonate.
6. Sodium Phosphate precipitates yellowish-white uranyl phosphate,

UOe2 (NOs) a+ NazgHP O4 —- 2NaNOsz +UO2HPO,,

while, in the presence of ammonium acetate, uranyl ammonium phos-
phate is precipitated:

NazgHPO, +U0e2 (NOsz)2 +NH4CoH302 ae
2NaNOz +HCe2H302 +U02 . NH,4POx,.

Both precipitates are insoluble in acetic acid, but soluble in mineral
acids.

7. Potassium Ferrocyanide produces a brown precipitate, or in
very dilute solutions, a brownish-red coloration,

Fe(CN)¢ ~+2U02t* — (UO2)2[Fe(CN)<],

On addition of potassium hydroxide the brownish-red_precipi- !
tate becomes yellow, owing to the formation of potassium uranate:

(UO2)2[Fe(CN)6]+6KOH = Ka[Fe(CN)¢6]+3H20+K2U207.

_ (Distinction from cupric ferrocyanide.)

REACTIONS IN THE DRY WAY

The borax (or sodium metaphosphate) bead is yellow in the oxidizing —
flame and green in the reducing flame.

TITANIUM 157

TITANIUM, Ti. At. Wt. 48.1
Sp. Gr. =4.87. M. Pt. =1900° (?)

Occurrence.—Titanium occurs in nature most frequently as the
dioxide, rutile (tetragonal), anatase (tetragonal), and brookite (ortho-
rhombic). Titanium is also found in the minerals perowskit, CaTiOs,
titanite, CaSiTiOs, and ilmenite, FeTiO3, as well as in many crystalline
rocks. It is present in most rocks, but usually only in very small
quantity.

Titanium itself is a gray metal, very similar to iron. On being
heated in the air it burns brightly to white titanium oxide. The
following oxides of titanium are known: TigQ2, TizO3, TiOe, TiO3 or
possibly TiO2-H2QOz.

The oxides TigO2 and TigO3 form violet-colored salts, which are
readily changed by oxidizing agents into derivatives of TiOz. The
most important oxide is titanium dioxide, which sometimes acts as
a base and sometimes as an acid. The titanium dioxide as it occurs
in nature (rutile, etc.) is insoluble in all acids. In order to bring it
into solution it is best to fuse it with potassium pyrosulfate, whereby
it is changed into titanium sulfate:

Ti02+2K28207 = Ti(SO4)2+2K280u.

The melt is dissolved in cold water. It can also be dissolved by fusing
with sodium carbonate and treating the melt with 6 N hydrochloric
acid.

REACTIONS IN THE WET WAY

For these reactions a solution of titanium sulfate or of titanium
hydroxide in hydrochloric acid may be used.

1. Potassium Hydroxide precipitates, in the cold, gelatinous
orthotitanic acid,

++
Tit++40H- > HuTiO,,

which is almost insoluble in an excess of the reagent, but readily soluble
in mineral acids.

If the precipitation by potassium hydroxide takes place from a
hot solution the titanium is precipitated as metdtitanic acid,

++
Tit ++40H — H20+H2TiOs,

which is difficultly soluble in dilute acids. By long digestion with
concentrated hydrochloric or sulfuric acid it goes gradually into
solution. By the ignition of both these titanic acids the anhydride

158 REACTIONS OF THE METALS

TiO>s is obtained, which is only slightly soluble in concentrated hydro-
chloric acid, but readily soluble in hot concentrated sulfuric acid.

2. Ammonia, Ammonium Sulfide, and Barium Carbonate (like
potassium hydroxide) precipitate, in the cold, orthotitanic acid, readily
soluble in acids; and from hot solutions the difficultly soluble meta-

titanic acid.
TiCl4+4N H40OH =4NH4Cl+ HygTiOg,

TiCla+ 2(NH4)2S+4H20 =4NH4Cl+2H2S 7 +HyTiO,,
TiCl,+2BaCO3+2H20 =2BaCle +2COs T +HuTiOs.

3. Alkali Acetates precipitate on boiling all of the titanium as
metatitanic acid:

++
Ta +44CoH302 +3H20 =4HC2H302+ HeTiOs3.

Titanium acetate is first formed, but it is completely decomposed
hydrolytically by boiling the dilute solution.

4. Water.—Not only titanium acetate is hydrolytically decom-
posed by water, but all titanium salts. |

The ease with which soluble titanium salts undergo hydrolysis with the
formation of insoluble titanic acid is the basis of the several methods for
separating titanium from aluminium, iron, chromium, ete.; the oxides of these
metals are fused with potassium pyrosulfate, the product of the fusion is
dissolved in cold water,* and the solution is then heated to boiling. The tita-
nium is completely precipitated as granular metatitanic acid, which can be’
readily filtered off, while the remaining metals remain in solution as sulfates:

Ti(SO,) +3H.0 = 2H.80,+-H.TiO;.

As this reaction, like all hydrolytic decompositions, is reversible, it is evident
that, in order to make the precipitation of the metatitanic acid complete, the
amount of free acid present should be kept small,{ considerable water should be
used, and the solution kept hot while filtering.

In order to precipitate titanic acid from a solution according to this method,
the concentrated solution it treated with sodium carbonate in the cold until
a slight permanent precipitate of Ti(OH), is obtained, sulfuric acid is added
drop by drop until the precipitate is just dissolved, considerable water is added
(300 to 500 cc. of water should be used for each 0.1 gm. TiO.) and the solution
kept at the boiling temperature for one hour. The granular metatitanic acid
thus obtained is easy to filter as long as free acid is present. On being washed
with pure water a turbid filtrate is always obtained; a little dilute sulphuric
acid should therefore be added to the wash-water.

* Solution takes place much more quickly if the liquid is kept in constant motion,
e.g., by conducting a current of air through it.
+ If too little acid is present, however, iron and aluminium will precipitate.

TITANIUM 159

With the separation of titanium according to the above-described method,
a play of colors will be observed in the bottom/of the glass beaker or flask which
is very characteristic of titanic acid.

The presence of tartaric acid, citric acid, and many other organic compounds
prevents the above reactions. ‘In such a case the organic substance must be
first destroyed either by ignition or by oxidation with potassium permanganate
(see pp. 129 and 152), the titanium dioxide dissolved in sulfuric acid and pre-
cipitated according to any of the above methods,

5. Potassium Ferrocyanide produces in slightly acid solutions
a brown precipitate.

6. Tannin produces a brown precipitate, which soon becomes
orange.

7. Sodium Thiosulfate precipitates in boiling solutions all of the
titanium as metatitanic acid:

TiCk,+ 2Na25203-+ H20 =4NaCl+2S+2802 T +HoeTiOs.
8. Sodium Phosphate precipitates basic titanium phosphate,

TiCl4a+3Na2HPO4s.+ H20 =4NaCl+2NaH2P04+Ti‘PO,) (OH),

soluble in mineral acids, insoluble in acetic acid.

9. Hydrogen Peroxide.—If hydrogen peroxide is added to a slightly
acid solution of titanium sulfate, the solution is colored orange red,
except in the presence of small amounts of titanium, when the color is
light yellow. This reaction, which depends upon the formation of
TiO3, or possibly TiO2-H2QOs2, is exceedingly delicate, and is especially
suitable for the detection of titanium in rocks. Vanadic acid behaves
similarly with hydrogen peroxide.

If a solution of titanium sulfate is treated with a large excess of
hydrogen peroxide and then potassium hydroxide is added, a precipitate
is formed which dissolves in a great excess of the alkali, forming a
yellow solution. This solution remains clear for a long time, but
eventually a bright yellow precipitate of Ti(OH)4 is formed.*

An insoluble titanium compound on being fused with sodium
peroxide in a nickel crucible yields a melt which permits the extraction
of all the titanium by water.+ If the solution is made strongly acid
with sulfuric acid, the orange-red color of pertitanic acid is apparent.
If iron was present in the original insoluble titanium mineral, it is left
insoluble in water after the fusion with sodium peroxide,

* A. CLAssEN, Ber., 21 (1888).
tJ. H. Watton, Jr., J. Am. Chem. Soc. (1907).

160 REACTIONS OF THE METALS
10. Zinc or Tin .produces in acid solutions, preferably hydro-
chloric acid, a violet color caused by the formation of TigCle:

2TiCl4 -}- He = 2HCl + TisCle.

The quadrivalent titanium compounds are not reduced by hydrogen
sulfide or sulfurous acid.

11. The Fluoride is quantitatively éhanged to the dioxide by
evaporation with sulfuric acid (difference from silicic acid).

TiF1+2H2S04=4HF T +2803 T +TiOo.

12. Ether does not dissolve titanium chloride. By shaking the
hydrochloric acid solution with ether, therefore, it is possible to separate
ferric chloride from titanium chloride.

13. Cupferron cf. (p. 152) precipitates the yellow titanium salt of
phenylnitrosohydroxylamine from acid solutions:

++
Tit *+4CsH;sNO-NO™~ —> Ti(CsHs;NO-NO)s.

In the presence of tartaric acid, iron can be removed by ammonia
and ammonium sulfide and the titanium precipitated in the acidified
filtrate with cupferron.

REACTIONS IN THE DRY WAY

Titanium compounds do not color the borax or sodium meta-
phosphate bead in the oxidizing flame; after continued heating in’
the reducing flame the bead becomes yellow while hot and violet
when cold. By the addition of a little tin the violet color appears
much more quickly. The addition of iron causes a brownish to red
bead.

On fusing titanic acid with sodium carbonate, sodium metatitanate
is formed, which is readily soluble in acids. By treatment with hot
water, sodium metatitanate is decomposed, forming metatitanic acid,
which is difficultly soluble in dilute acids. Even in cold water hydrolysis
of sodium titanate takes place and none of the titanium dissolves,

MANGANESE, Mn. At. Wt. 54.93
Sp. Gr. =about 8.0. M. Pt. =1225° C.

Occurrence.—The most important manganese minerals are pyrolu-
site, MnOzg, orthorhombic; polianite, also MnOog, tetragonal, isomor-
phous with rutile and tinstone; braunite, Mn2Os, tetragonal; man-
ganite, HMnOz, orthorhombic, isomorphous with géthite and diaspore;
hausmannite, MngQO4, tetragonal; and rhodochrosite, MnCO3s. Man-

* MANGANESE 161

ganese is a constant companion of iron, so that we find it in varying
amounts in almost all iron ores.

It is ‘a grayish-white metal which is readily oxidized in moist air,
and is attacked by dilute acids, even acetic acid. It forms the follow-
ing oxides:

MnO, Mn2Qs, Mns3QO0u4, MnOz, (MnQOs), Mne207.

By treating any one of these oxides, except MnO, with cold, dilute

hydrochloric acid a dark, greenish-brown solution is obtained, which

on. being heated evolves chlorine and becomes colorless. The solution
then contains a bivalent manganese salt,—a derivative of MnO.

Mn0O+2HCl=H20+MnCle,

Mn203 +6HCl=3H20 +2MnCle+Cle T ,
Mn304 +8HCl=4H20 +3MnCly+Cle T ,
Mn0O2+4HCl=2H20 +MnCly+Clo T ,
Mn:207 +14HCl=7H20 +2MnCle+5Cle 7.

All manganese oxides dissolve on warming with concentrated
sulfuric acid, forming manganous sulfate, accompanied (except with
MnO) by evolution of oxygen:

MnO +H2S04=H20 +MnS80,,

2Mn203 + 4H2SO4 =4H20 +4Mn804+02 Tf,
2Mn304 +6H2801 =6H20 +6MnS04+02 T,
2MnOz +2H 2804 =2H20 +2MnS04+02 1.

The behavior of the higher oxides, MnOz, MneQ3, and Mn3Q0q,
with boiling dilute nitric or sulfuric acid is very interesting: MnOs
is not attacked at all by these dilute acids; while Mn2Oz3 gives up
half of its manganese to the acid, the other half remaining undissolved
as brown hydrated manganese dioxide, H2MnOs; two-thirds of Mn304
is dissolved by these acids, brown hydrated manganese dioxide being
left behind, as before.

This HeMnO3 separates out just as metasilicic acid is deposited
from a silicate on the addition of a strong acid:

CaSiOz + QHt > HeSi03 + Cat*.

In fact, hydrated manganese dioxide behaves in most cases exactly
like an acid, the oxides Mn2Oz and MngQOx behave like manganous salts
of this acid and are to be regarded as manganites.

162 REACTIONS OF THE METALS

Mn203 therefore is to be regarded as manganous manganite,
O Ne
Mn 0 MnO, of analogous composition to manganous carbonate,

MnCOs, and manganous metasilicate, MnSiO3.

According to this conception, electrolytic dissociation shou!d give
rise to Mn** cations and MnO; anions; it is easy to understand, there-
fore, why Mne2O3 gives up half of its manganese on treatment with
dilute nitric acid, with the separation of manganous acid:

Mn:-Mn0O;+2HNO;3 @ Mn(NOs3)2+H2MnOs.

Mn3zQ4, which gives up two-thirds of its manganese, may be con-
sidered to be the manganous salt of orthomanganous acid, Hy4MnQ,.

On treating Mnz-MnO, with nitric acid, the ortho acid first sep-
arates out; it loses water, and goes over into metamanganous acid:

Mn2-Mn0,+4HNO;3 = 2Mn(NOs3) 2+H4Mn0Oxz,
H4zMn0O, == HO +H2MnOs3.

MnOz stands in the same relation to HoMnOs as COz to HeCOs,
as SiO2 to HeSiOs, and as SnOzg to HeSnOs; MnOsz, therefore, behaves
like an acid anhydride.

Like SnO2g (which see), manganese dioxide behaves partly as an
acid anhydride and partly as the anhydride of a base, because it prob-
ably forms the chloride MnCl4. For if MnOz is treated with cold
concentrated hydrochloric acid it dissolves with a brownish-green
color, forming manganese tetrachloride, soluble in ether with a green
color. If, therefore, the aqueous solution of MnCl, is shaken with
ether, the upper layer is colored green.

Mn(S0,)2 and Mn,.(SO,); are also known. These salts are hydrolyzed
readily, forming H,2MnO; and Mn-Mn0O;:

Mn (SOx)2 +3H,0 = H,Mn0O; +2H.SO,,
Mn,.(SO,);+3H.0 @ Mn,0;+3H.S8O..

Not only manganous manganites are known, but quite a number
of other manganites. Some of these play a very important part in
analytical chemistry; as, for example, zine and calcium bimanganites,
ZnHe(MnOs)2, CaH2(MnOs)2, which are analogous to calcium bicar-
bonate, CaH2(COs)e.

Zinc bimanganite is formed in the volumetric determination of
manganese (see Vol. II). Calcium bimanganite is of importance.
technically. Thus the recovery of manganese, by the Weldon process,

MANGANESE 163

in the manufacture of chlorine, depends upon the formation of calcium
bimanganite.

Manganous oxide, MnO, is the only oxide of manganese which
in all cases acts as the anhydride of a base. By dissolving this oxide
in acids, manganous salts are always obtained, in which the manganese
is bivalent. ‘The oxide MnQOg has never been isolated, but there are
salts (ReMnOuz, see p. 169) known which are derived from it. Mng207
is a distinct acid anhydride, from which the permanganates (RMnQ,)

are derived.
In the study of the reactions of manganese we will consider first
the manganous compounds, then the manganates and permanganates.

A. Manganous Compounds

The manganous compounds are pink both in the crystalline state
and in aqueous solution; but in the anhydrous state they are colorless
with the exception of the sulphide.

REACTIONS IN THE WET WAY

1. Potassium or Sodium Hydroxide precipitates white manganous

hydroxide,
Mn*tt+20H- — Mn(OH)p,

which rapidly becomes brown in the air, owing to the formation of
manganous manganites, which are less soluble than Mn(OH)e.
First a part of the manganous hydroxide is oxidized by the air

- to manganous acid:
Mn(OH)2+0 — H2Mn0Os,

which, on coming in contact with the basic manganous hydroxide,
immediately forms a salt with it—a manganite,
H2Mn0O3+Mn(OH)e2 <2 Mn-Mn0O3+2H20
or possibly,
2HeMnO3 +Mn(OH)s2 = MnHo2(Mn0Osz)e +2H>0.
This oxidation takes place in the air only gradually, but imme-

diately in the presence of chlorine, bromine, hypochlorites, hydrogen
peroxide, etc.:

Mn(OH)2+2Na0H+ Cle = 2NaCl+ H20 +H2MnOs,
Mn(OH)2+H202 = H20+H2Mn0O3.

164 REACTIONS OF THE METALS

The formation of manganites is of technical importance, as mentioned on
p. 162. The residue obtained in the preparation of chlorine from pyrolusite
and hydrochloric acid consists chiefly of manganese chloride; by adding lime
to it, manganous hydroxide is formed. This mixture of manganous hydroxide
and lime is exposed to the action of the atmosphere, whereby manganous acid
is formed, which unites with the calcium as the stronger base, forming calcium
bimanganite, so that finally all of the manganese is oxidized to manganous acid:

2Mn(OH), +0, +Ca(OH), = 2H.0 +CaH,(Mn0O;)>.

On treating the residue, when in the right condition. with hydrochloric
acid again, the same amount of chlorine is obtained as from the original pyro-
lusite:

2Mn0O.+S8HCl = 4H,.0+2MnCl, +2Cl, T 4

CaH.(MnO;),+10HCl =6H,0+2MnCl,+CaCl,+2Cl, T.

It is, however, necessary to add a little more hydrochloric acid in the latter
case, because a part of the acid is used up in setting the An gSEos acid free
from the manganite.

2. Ammonia precipitates (as with magnesium and ferrous salts)
from neutral solutions free from ammonium salts a part of the man--
ganese as the white hydroxide:

MnCle+NH40OH @ Mn(OH)2+2NH4Cl.

If sufficient ammonium chloride is present, ammonia causes no
precipitation (cf. p. 94). The greater part of the manganese then
remains in solution as manganous chloride, but a small amount exists as
the hydroxide. On standing in the air, this dissolved hydroxide is
changed slowly. into the more difficultly soluble manganous acid, which is
deposited in brown flocks. The condition of equilibrium in the solution
is thereby disturbed, and in order to restore it more hydroxide is formed,
and the reaction continues in this way until finally all of the manganese
may be precipitated. This fact must be considered in the separation
of manganese from ferric iron, aluminium, ete. If a solution of ferric
and manganous chlorides contains sufficient ammonium chloride, none
of the manganese and all of the iron will be precipitated on the addition
of ammonia, but, if the solution stands in contact with the air, little
by little the manganese will be precipitated. In effecting the separa-
tion, therefore, an excess of ammonium chloride should be present, the
solution boiled to remove the air as much as possible from the solution,
then a slight excess of ammonia should be added and the solution fil-
tered immediately. The separation even then is not quantitative, but
is satisfactory for qualitative analysis.

Oxidizing agents in the presence of ammonia cause the precipitation of
manganese as H,MnO;. Bromine is ordinarily used as the oxidizing agent, but

MANGANESE 165

a number of precautions are necessary to accomplish the complete precipitation
of manganese by means of bromine and ammonia.

If a neutral solution of manganous salt is treated with bromine, the precipita-
tion of manganese as H:Mn0Q; is always incomplete:

Mnt +4 Br.+3H.O y ax 4 4Ht++2Br-+H.Mn0O,.

The precipitation of the manganese can be made complete, in accordance with
the mass action principle (p. 13) if the hydrogen ions formed in the reaction
are neutralized; sodium bromide in neutral solution will not reduce H.MnO,,
but hydrobromic acid will do so. The solution may be neutralized by caustic

alkali, alkali carbonate, alkali acetate (cf. p. 45), or ammonia. A solution of

manganous salt may contain a considerable excess of acetic acid and yet the
manganese will be completely precipitated by bromide in the presence of
sodium acetate.

Ammonia is not altogether satisfactory as a neutralizing agent in this case
because it reacts with bromine as well as with hydrobromic acid. When
bromine is added to ammonia solution, a vigorous reaction takes place and
nitrogen is evolved (cf. p, 92);

8NH,OH+3Br, =6NH,Br+8H,.0+N, t

When bromine is added to a solution of ammonium chloride a very slow oxida-:
tion results and nitrogen gas is evolved little by little:

2NH,*+3Br, @ 8Ht+6Br +N; fT.

If sodium acetate is added to this solution, the reaction is accelerated greatly as
a result of diminishing the concentration of the hydrogen ions and, by heating
the solution, which increases the speed of the reaction and causes the rapid
expulsion of nitrogen gas, all of the ammonium salt can be decomposed by
means of a very slight excess of bromine.

The presence of ammonium salts prevents the precipitation of manganese by
bromine and sodium acetate; when all the ammonium salt has been oxidized to
nitrogen, the manganese can be precipitated as H,MnQ;.

The conditions are more favorable for the precipitation of manganese when
the solution contains hydroxyl ions, as in an ammoniacal solution. The addi-
tion of bromine to such a solution usually results in the immediate precipitation
of some of the manganese but, as a result of the action of bromine on ammonia
or ammonium salt, the solution usually becomes acid and the precipitation of
the manganese is then incomplete. To precipitate all of the manganese by
means of ammonia and bromine, it is best to proceed as follows:

Dilute the solution to about 200 cc. and neutralize, if necessary, with am-
monia. Add a little bromine water, a slight excess of ammonia and stir to
promote the formation of a precipitate. Heat the solution, add a little more
bromine water and make slightly ammoniacal again. In case the amount of the
manganese precipitate is perceptibly increased by this last treatment with
bromine and ammonia, repeat the operation. Filter off the precipitated
H:Mn0O;, concentrate the solution somewhat by evaporation and again treat
with bromine water and ammonia. Sometimes H,MnO, is precipitated on the
sides of the vessel during evaporation. The treatment with bromine and

166 REACTIONS OF THE METALS

ammonia should be continued until a filtrate is obtained which will not give
any more precipitate with these reagents, *

3. Alkali Carbonates precipitate white manganous carbonate,
Mntt*+C03" > MnCOs,

which after long boiling is changed by the oxygen of the air into less-
soluble, hydrated manganese dioxide: :

2MnC03+2H20+02=2CO2 | +2H2MnOs.

4. Ammonium Carbonate precipitates even in the presence of
ammonium salts the white carbonate (difference from magnesium).
5. Barium Carbonate produces a precipitate only in hot solutions.
6. Sodium Phosphate precipitates white, tertiary manganous
phosphate,
4HPO4.-+3M nits 2H2PO04-+Mnsz(POx4)2,

soluble in mineral acids and in acetic acid:
Mn3(PO.4)2+2H*t @ 3Mn*t+2HPO.-.

If to the boiling solution of this precipitate in acid an excess of
ammonia is added, manganous ammonium phosphate will be precipi-
tated, as with magnesium (see p. 95):

| HPO.4z-+NH40H @ NH4t+ PO4=+ H20.
Mntt-+NH,t+ PO4=+7H20 als Mn(NH,4)PO,4 . 7H2O.

The precipitate consists of pink scales and is practically insoluble in
water.

7. Lead Peroxide and Concentrated Nitric Acid. (Volhard’s
reaction).—If a solution containing only traces of manganese is boiled
with lead peroxide and concentrated nitric acid, then diluted with
water and the residue allowed to settle, the supernatant liquid acquires
a distinct violet-red color, owing to the formation of permanganic acid:

2Mn*+t-+5PbO2+6H*t > 5Pbt*+2H20+2HMn0Og.

This extremely delicate reaction does not take place in the presence
of much hydrochloric acid or chlorides, because the permanganiec acid
is thereby destroyed:.

2HMn0O4+14HCl=8H20+ MnCle+5Cle T .

* The above explanation is given at length because of its importance in quantita-
tive analysis. The facts upon which the explanation is based have been carefully
verified by quantitative experiments performed in the laboratory of the translator.

MANGANESE 167

8. Sodium Bismuthate added to a cold solution of a manganous
salt in dilute nitric acid (about sp.gr. 1.13) causes the formation of
permanganic acid. The reagent, which corresponds approximately to
the symbol NaBiOs, is prepared by fusing bismuth oxide with sodium
peroxide; it is insoluble in water and the excess of reagent may be
filtered off through asbestos after applying the test. The reaction may
be expressed by the equation:

2Mnt*t+5NaBi03+16Ht > 5Nat+5Bitt++7H2,0+2HMn0,.

The test is extremely delicate when nothing is present that will
react with the permanganate formed. An insoluble carbonaceous
‘residue, such as remains after the solution of cast iron in acid, must be
filtered off before adding the reagent. If the solution is heated, the
permanganic acid breaks down and hydrated manganese dioxide is
precipitated.

Bismuth dioxide, BiOz2, may be used instead of sodium bismuthate.

9. Ammonium Persulfate. If a hot solution of a manganous
salt in either dilute sulfuric or nitric acid is treated with ammonium
persulfate, all the manganese is gradually oxidized to the quadrivalent
condition and a precipitate of hydrated manganese dioxide is formed:

Mn*t+820.7+3H20 — 2804-+4H*+ HeoMnOs.

If, however, the solution contains a trace of silver nitrate as catalyzer,
then the oxidation goes farther and permanganic acid is formed:

2Mn*t-+5S820s-+8H20 — 16Ht+10S8S04-+2Mn0.-.

This reaction is quantitative for small amounts of manganese and
in the absence of anything that will react with the permanganic acid.*

10. Ammonium Sulfide precipitates from manganese solutions
flesh-colored, hydrated manganese sulfide:

Mn**+S~— MnS.

On boiling with a large excess of ammonium sulfide it is changed
into less hydrated green manganese sulfide of the formula, 3MnS+ H20.

The solubility product of MnS is relatively large (cf. p. 22) and to
precipitate all the manganese as sulfide an excess of S~ ions is necessary.
The precipitate dissolves readily in dilute acid as a result of the removal
of the S~ ions to form non-ionized hydrogen sulfide (cf. p. 47).

11. Potassium Cyanide.—On adding potassium cyanide to a solu-
tion of a manganous salt, a brown precipitate appears which dissolves

* Cf. M. MarsHat, Z. anal. Chem., 48 (1904), 418, 655.

168 REACTIONS OF THE METALS

in an excess of potassium cyanide, forming a brown solution. On
standing, or by heating the solution, a voluminous green precipitate of
K[Mn(CN)3] is formed which is soluble in strong potassium cyanide
solution:

Mn**+2CN- @ Mn(CN)2, Mn(CN)2+KCN = K[Mn(CN)3]
K[Mn(CN)3]+3CN~ = K*+[Mn(CN)6¢]-~.

To keep the manganese in solution in the form of Mn(CN)¢ ~ ions,
it is necessary to use an excess of potassium cyanide. If the concen-
tration of the cyanide is diminished by dilution, some green K{Mn(CN)s]
is formed and if the dilute cyanide solution is boiled, a precipitate of
Mn(OH): results:

Mn(CN)6 ~+2H20 @ 2HCN+4CN +Mn(OH)>.

The stability of these complex cyanides, therefore, is much less than that
of the corresponding nickel compounds (p. 175). -This permits an
interesting method of separating nickel from manganese.

Nickel sulfide is much less soluble than manganous sulfide (cf.
p. 22) so that it is possible to precipitate nickel as sulfide in the presence
of acetic acid and sodium acetate; under these conditions no manganese
sulfide is formed. If, on the other hand, ammonium sulfide is added to —
a hot, dilute solution containing the complex cyanides of nickel and
manganese, the nickel will remain in solution and the precipitation of
the manganese as sulfide will be complete:

Mn(CN)¢ +S <6CN +MnS.

In the presence of a large excess of potassium cyanide, however, the
Mn(CN)¢5 ~ anion is so stable that none of the manganese is precipi-
tated in the cold by ammonium sulfide.

12. Potassium Chlorate. By boiling a solution of manganous salt
in concentrated nitric acid with an excess of potassium chlorate, all
of the manganese is precipitated as MnQz.

Mn(NO3)2+2KCl1O3 ~Mn0O2+2KNO03+2Cl0sz.

REACTIONS IN THE DRY WAY

The bead of borax, or salt of phosphorus, is amethyst. red after heat-
ing in the oxidizing flame with small amounts of manganese, almost
brown with larger amounts, and can then be mistaken for the nickel
bead. Heated in the reducing flame, the manganese bead becomes
colorless, while the nickel bead appears gray.

to mee oe

MANGANESE 169

On fusing any manganous compound with caustic alkali or alkali
carbonate (on platinum foil) in the air, or, better still, in the presence
of an oxidizing agent (such as potassium nitrate, potassium chlorate,
etc.), a green melt is obtained, owing to the formation of the alkali
salt of manganic acid, as is shown by the following equations:

Mn0O+Na2C03+02=CO2 | +NazMnO,,

Mn0O2+ NuzC03+0 =CO2 | +Na2MnOa,
Mn203+2Naz2C03+30 =2CO2 7 +2Na2MnOz,
Mn304+3Na2CO03+50 =3CQO2 | +3Na2MnOz,
MnS04+2Na2C03+ O02 =2COz2 | +NaeSO4+ NazMnO,.

The oxygen comes either from the air or from the nitrate or chlorate:
KNOz3=KNO2+0, KClO3=KCI+30.

This reaction is exceedingly delicate; a fraction of a milligram
of any manganese compound can be recognized by the formation of

this green color.
By ignition in the air the oxides of manganese are changed to Mn30a:

3Mn0+0 = Mng0Quz, 3MnO2 = Mnsg04+O02 T ;
6Mne203 =4Mn304+ O2 T ‘

B. Manganic and Permanganic Acids

The free manganic acid has never been isolated. If we attempt
to form it from the green melt of the alkali manganate by the addition
of acid, permanganic acid and hydrated manganese dioxide will be
obtained; a part of the unstable manganic acid oxidizes another part
of the same to permanganic acid, while the oxidizing part is itself
reduced to hydrated manganese dioxide:

3HeMnO4 = 2HMn0O4+HeMnO03+ H20.

This transformation takes place so readily that the green solution
of the manganate is changed to a reddish-violet solution of a permangan-
ate by simply standing in the air, with the help of the carbonic acid
which the air always contains:

3KeMn04+2C0e+ H20 = 2KeCO03+ H2Mn0O3 +2KMn0Osg.

170 REACTIONS OF THE METALS

;

The reaction takes place much more rapidly, however, if a few
drops of a strong acid are added.

The oxidation of one molecule at the expense of another of the same
kind is of quite common occurrence in chemistry. It always involves a
loss in the available or free energy which the molecules originally possessed.
The total energy possessed by any molecule can be considered to consist partly
of free energy and partly of unavailable energy. A reaction that takes place
spontaneously is always characterized by the fact that the free energy of the
system is less afterward than it was before the reaction took place. The coi-
dition with the smallest free energy is the most stable condition.

It might be inferred that the most stable conditions are those having the
smallest quantities of energy, but a little consideration shows that in promoting
chemical reactions it is not so much the total energy as it is the available energy
which comes into consideration. Thus, there is a vast amount of energy stored
up in the heat of the ocean, but it is not available energy, because it is in sur-
roundings at the same temperature. The air under ordinary atmospheric pres-
sure could perform a great deal of work if it were brought in contact with a
space in which a much lower gas-pressure prevailed, but otherwise the vast
amount of energy is not available.

In the changes that take place with any given element it is not necessarily
true that an increase in the total energy will always involve an increase in the
free energy associated with the element. The fact that the solution of a man-
ganate, in which the valence of the manganese is six, decomposes readily
indicates merely that the free energy in the system composed of quadrivalent
and heptavalent manganese is less than the free energy involved in the system
containing all the manganese in the hexavalent condition. Frequently the
conditions are just the reverse and the most stable condition is one of inter-
mediate valence. Thus permanganate and manganous salt react to form quad-
rivalent manganese. A few reactions similar to the decomposition of HzMnO,
will be given.

Hypochlorites are changed, by warming the aqueous solution, into chlorate
and chloride; one atom of chlorine is oxidized from the valence of one to the
valence of five at the expense of two atoms of chlorine, which are reduced from
a positive valence of one to a negative valence of one;

3NaClO =NaClO;+2NaCl.

Ignition of a chlorate causes the formation of a perchlorate, a chloride and
free oxygen. Here one atom of chlorine is increased two in valence, oné
atom of chlorine loses six charges and the remaining four charges cause the
oxidation of two atoms of negatively charged oxygen:

2NaClO; = NaClO,+NaCl+0, fT.

Nitrous acid is changed in aqueous solution into nitric acid and nitric oxide, one
atom of nitrogen gaining two charges and two similar atoms losing one
charge:

MANGANESE 171

Hypophosphorous acid and also phosphorous acid can be changed into Boe:
phoric acid and phosphine:

2H;PO, =H,;P0O,+PH; f , 4H;PO, =3H;P0,+PH; T .

Alkali thiosulfates and alkali sulfites are changed by ignition into sulfate
and sulfide:

4Na.8.0; = 3Na.SO.,+ NaS; ’ 4Na.SO; = 3Na.S80,+ NaS.

Permanganic Acid, HMnO,, although much more stable than
- manganic acid, is known only in aqueous solution; but the anhy-
dride Mn207 has been isolated. On cautiously adding concentrated
sulfuric acid to the cooled solution of a permanganate, oily drops of
reddish-brown Mn207 separate out, which, however, on being warmed
(the heat of reaction ‘is sufficient), explode with scintillation:

2Mn207 +4H22SO4 =4MnS04+4H20+502 t .

The salts of permanganic acid (the permanganates) are all soluble
in water, with a reddish-violet color, and are very energetic oxidizing
agents. In acid solution the heptavalent manganese in permanganate
is usually reduced to bivalent manganous salt, but in alkaline, or nearly
neutral, solutions manganese dioxide is the usual product.

Oxidation in Acid Solution. Typical oxidation equations with
permanganate have already been explained on p. 31. <A few of
these reactions will be repeated here, but it will not be necessary to
enter into further details concerning the method of balancing the equa-
tions. To avoid exact repetition, the equations will be given with the
entire molecules written instead of merely the ions involved,

2KMn04+3H2804+ 10HCl = K2804-+2MnS804+8H20+5Cle J ,
2KMn04+3H2504+ 10HI = K2804-+-2MnS04+8H20 +52,
2KMn04+3H2804+5He2S = Ke8O4+2MnS04+8H20-+5S8,
8KMn04+12H2S04+5PH3 =4K2804+8MnS04+ 12H20+5H3PO4.
2KMn04+6802+2H20 = 2KHS04+2MnS04+ Hed20¢.

In this last equation the proportion of sulfate and of dithionic acid,
H2S820¢, will vary with the temperature and concentration of the solu-
tion:
2K Mn04+5H2eC204+3H2S04 = Ko804+2MnS04+8H20+ 10C02 Lie
2KMn04+5H202+4H2S04 = 2KHS01+2MnS01+8H20+502 7 ,
2KMn04+5KeC2064+ 14H2S04=

= 12KHS04+2MnS0.i+8H20+ 10CO02+502 7.

172 REACTIONS OF THE METALS

Persulfuric acid, which is analogous to percarbonic acid, does not
reduce a solution of a permanganate. An interesting reaction is
that which takes place in nearly neutral solution between permanganate
and manganous ions. The principal product is MnOz, which will carry
down some of the bivalent manganese as manganous manganite unless
an excess of some other ion is present which forms an insoluble man-
ganite, e.g., zinc or calcium ions;

2KMn0O4 +3MnSO4 + 2HeO = 2KHSO,4 -+- 5MnO> +H2S0..

Oxidation in Alkaline Solution—Many organic substances are
oxidized by permanganates in alkaline solution with precipitation
of manganese dioxide. Thus formic acid is oxidized to carbonic
acid, ethyl alcohol to aldehyde and acetic acid, cellulose (paper) chiefly
to oxalic acid, so that a. solution of a permanganate cannot be
filtered through paper. By boiling a concentrated solution of potassium
permanganate with concentrated potassium hydroxide, potassium
manganate is formed with evolution of oxygen, and the color of the
solution becomes green:

4KMn04--4KOH = 4KoMn04-+2H20+0s> 1.

By heating solid potassium permanganate to 240° C., potassium
manganate is formed, also with evolution of oxygen:

2KMn04= KeMn0O4+Mn0Oe +O2 T .

NICKEL, Ni. At. Wt. 58.68

Sp. Gr. =8.9. M. Pt. =1452°C.

Occurrence.—In the native state nickel occurs only in meteor-
ites. It is most frequently found in combination with sulfur, arsenic,
and antimony in regular and hexagonal crystallizing minerals, of which
the following are the most important:

Isometric System: Chloanthite, NiAse; gersdorffite, NiAsS; ullman-
. nite, NiSbS. Hezagonal System: Niccolite, NigAse; breithauptite,
NigSbe2; millerite, NigSe.

Nickel also occurs as regular crystals of bunsenite, NiO, isomor-
phous with periclasite, MgO, and manganosite, MnO; as garnierite
or noumeite, He(NiMg)SiO4-+aq, a mineral occurring in New Cale-
donia, from which pure nickel can be prepared; and finally as anna-
bergite, Nig(AsO4)2-8H2O, isomorphous with erythrite.

NICKEL | 173

Metallic nickel possesses a silver-white color and is difficultly
soluble in hydrochloric and sulfuric acids, but readily soluble in nitric
acid. It forms two oxides: green nickelous oxide, NiO, and brownish

black nickelic oxide, NigOs.
By dissolving either of these oxides in acids, salts of bivalent nickel
are always obtained:

Ni0+2HCI=H20+NiCh,
Ni.03+6HC1=3H20+2NiCle+Cl ft ,
INi203-+4H2804 = 4H20 + 4NiSO2+02 1 .

Nickelous oxide behaves as a basic anhydride, but nickelic oxide
acts as a peroxide and forms no salts.

The crystallized salts of nickel and their aqueous solutions are
green, but in the anhydrous condition they are usually yellow. Most
of the salts are soluble in water; the sulfide, carbonate, and phosphate
are insoluble.

REACTIONS IN THE WET WAY

1. Potassium Hydroxide precipitates apple-green nickelous hydrox-
ide,
Nitt+20H~— Ni(OH)s,

insoluble in excess of the precipitant, readily soluble in acids.
2. Ammonia precipitates (in neutral solutions free from ammo-
nium salts) a green basic salt,

2NiSO4+2N H40H = (NH4)2804+NieSO,4- (OH)2,

soluble (with a blue color) in excess of ammonia, forming complex
nickel ammonia ions (cf. p. 25).

Ni2SO.- (OH)2+12NH3 — 2Ni(NH3)6"*+20H~+S04°-.

In the presence of sufficient ammonium salt, ammonia produces
no precipitate, as with magnesium, ferrous and manganous salts; potas-
sium and sodium hydroxides, however, precipitate the green hydroxide
(difference from cobalt, see p. 179).

The anhydrous chloride and sulfate readily absorb ammonia,
forming anhydrous nickel ammonium salts:

NiCle+6NH3 =[Ni(NHs3)6]Cl2; NiSO4+6NH3 =[Ni(NHs3)6|SOx.

174 REACTIONS OF THE METALS

The nickel ammonia cations are very stable in the presence of
an excess of ammonia. In pure water they are in equilibrium with
a small quantity of nickel ions of dissolved ammonia, and ammonium
hydroxide: :

Ni(NH3)6** — Ni**+6NH3; NH3+H20 — NH4OH.

3. Potassium and Sodium Carbonates precipitate apple-green
nickel carbonate:
Ni**+CO37 > NiCOs.

4. Ammonium Carbonate behaves similarly, but the precipi-
tate which is formed is soluble in an excess of the precipitant, form-
ing nickel ammonia carbonate.

5. Sodium Hypochlorite precipitates in the presence of alkalies
all of the nickel as brownish-black nickelic hydroxide, Ni(OH)s3.
Nickelous hydroxide is first formed by the alkali present, but it is
then oxidized by the hypochlorite to nickelic hydroxide:

2Nit*+40H-+ClO~+H20 — 2Ni(OH)3+Cl-.

On adding chlorine or bromine to the nickel solution to which alkali
has been added, nickelic hydroxide is likewise formed:

2Ni(OH)2+20H +Cle — 2Cl"-+2Ni(OH)s.

6. Barium Carbonate produces in the cold no precipitation; but*by
continued boiling, all of the nickel is thrown down as basic carbonate.

7. Hydrogen Sulfide precipitates no nickel from solutions which
contain mineral acid or much acetic acid; but from solutions slightly
acid with acetic acid and containing an alkali acetate, all the nickel is
precipitated as the black sulfide:

Nit *+2C2H302-+H2S — 2HC2H302+NiS.

.8. Ammonium Sulfide precipitates from neutral solutions the
nickel as sulfide:

NiCl2+(NH4)28 = 2NH4CI+NiS.

Nickel sulfide has a marked tendency to form colloidal solutions
of a dark-brown color, especially in the presence of ammonia or a con-
siderable excess of ammonium sulfide. By making the brown solution
slightly acid with acetic acid and boiling, the hydrosole is coagulated
and can be removed by filtration. The presence of ammonium salts
also favors the coagulation of the hydrosole,

NICKEL 175

If it is desired to precipitate the nickel as sulfide from an ammoniacal
solution, it is best to make the solution very slightly acid, add a little ammonium
chloride unless considerable is already present, heat to boiling, add colorless
ammonium sulfide drop by drop until no further precipitation takes place,
and then add 0.5 to 1 cc. of the reagent in excess. The nickel sulfide thus
obtained can be filtered without difficulty and the filtrate is free from nickel.
During the filtration care should be taken to keep the filter well filled with liquid
to prevent the oxidation of the precipitate, which takes place readily on exposure
to the air. To wash the precipitate it is well to use a hot, 5 to 10 per cent
ammonium chloride solution to which a little colorless ammonium sulfide
has been added. The washing can also be effected with hydrogen sulfide »
water without there being any danger of hydrosol formation.

Nickel sulfide is difficultly soluble in dilute mineral acids, readily soluble,
however, in strong nitric acid or in aqua regia, with separation of sulfur:

38NiS+6HCIl+2HNO; =3NiCl.+2NO T +4H,0+8S.

The sulfur usually separates out as a black film. This is caused by the
sulfur first melting, owing to the heat of reaction, enclosing small particles of
the black sulfide and protecting them from the action of the acid. By continued
action of the acid all the sulfide is dissolved, and the sulfur remains as yellow
drops, which are oxidized little by little to sulfuric acid,

S+2HNO, =H,.S0,+2NO tT e

Nickel and cobalt sulfides, though not precipitated by hydrogen sulfide
from a dilute, hydrochloric acid solution, dissolve with difficulty in a much
stronger acid. This is perhaps due to the fact that these sulfides exist in two
allotropic forms of different solubilities. The sulfide first precipitated is readily
soluble in acid, but on standing it becomes changed into a much more insoluble
condition. Most schemes of qualitative analysis are based upon this behavior;
the nickel and cobalt are separated from zinc and manganese by treating the
ammonium sulfide precipitate with cold, dilute hydrochloric acid. In such cases
some nickel and cobalt always passes into solution and the quantity dissolved
may be much larger than is ordinarily assumed. If the surface exposed to the
action of the acid is large, or if left in a finely divided state by the dissolving out
of other sulfide, a considerable quantity of nickel passes into solution in a com-
paratively short time. The reverse reaction, the precipitation of nickel sul-
fide by hydrogen sulfide in very dilute acid solution, also takes place very
slowly but continuously.*

10. Potassium Cyanide produces a light-green precipitate of
nickelous cyanide readily soluble in an excess of the precipitant, forming
potassium nickelocyanide:

Nitt+2CN- — Ni(CN)2; Ni(CN)2+2CN~ — [Ni(CN)a]~.
The Ni(CN)4~ anion is a stable complex, but it is decomposed by

* Cf. Noyrs, Bray and Spear, J. Am. Chem. Soc., 1908.

176 REACTIONS OF THE METALS.

the addition of acid. This is because He[Ni(CN),] is a very weak acid
and, like carbonic acid, is unstabie: |

Ni(CN)47-+2H* = He[Ni(CN)4]; He[Ni(CN)4] @ 2HCN+Ni(CN)o.

Hydrogen Sulfide in an Alkaline Solution containing a tartrate
gives a clear brown solution. This is a very characteristic test for
nickel and it enables the detection of as little as 0.2 mg. of nickel in
20 cc. of solution; the exact condition of the nickel in this solution is
not known. The main function of the tartrate is to form a complex
anion with nickel and thus prevent the precipitation of nickel hydrox-
ide. The characteristic brown solution is not obtained until the solu-
tion is nearly saturated with hydrogen sulfide. |

In alkaline solutions, containing an excess of cyanide, the [Ni(CN)4]~
anion is not dissociated to a sufficient extent into simple Nit* cations
to give a precipitate with ammonium sulfide (difference from man-
ganese and zinc) but it is readily decomposed by an oxidizing agent
such as chlorine, bromine, hypochlorite or hypobromite

QNi(CN)4]-+40H-+9BrO-+H20 — 2Ni(OH)3+8CNO~+9Br-,
or

QINi(CN)4]-+60H~+9Clz — 2Ni(OH)3+8CNCI+10CI-.

In the above reactions, first of all any excess alkali cyanide present is Oxi-
dized to cyanate or to CNCl. In the absence of an excess of cyanide ions, the
[Ni(CN),]~ begins to dissociate, [Ni(CN),J- ~ N iY +4CN~ , and as fast as the
ions are formed, they both become oxidized.

When cobalt i ions are treated with an excess of cyanide ions a cobaltocyanide
anion, [Co(CN) 4), is formed which is readily oxidized even by the oxygen
of the air to form cobalticyanide ions, Co(CN),>~. These anions are so stable
that they are not decomposed by treatment with oxidizing agents in alkaline
solution so that no cobalt is precipitated with Ni(OH); by the above treatment.
It is possible to detect 0. 2 mg. nickel in the presence of 300 mg. of cobalt by this
reaction.

The test can be applied to the solution obtained by dissolving the sulfides
of nickel and cobalt in aqua regia. The test can be made as follows: Evaporate
the solution nearly to dryness to expel most of the acid, add about 5 ee. of water —
and then sodium hydroxide solution, drop by drop, until the solution is
neutral or until a permanent precipitate is produced. Add potassium cyanide
solution, a few drops at a time, until all or nearly all of any precipitated cyanide
redissolves. Then add 0.5 to 3 ec. more of potassium cyanide solution accord-
ing to the probable amount of nickel and cobalt present. Heat to 50° or 60°
in an open dish for five minutes, or longer if the solution has not become light-
colored. This serves to oxidize the cobaltocyanide to cobalticyanide. Filter
off and reject any small precipitate that may remain. Add about 3 cc. of 2N
sodium hydroxide solution and conduct chlorine into the solution in the cold or
add a little bromine water. A precipitate of Ni(OH); should form within five
minutes,

NICKEL 177

11. Sodium Phosphate precipitates apple-green nickel phosphate,
3Nit +4 4HPO,= — 2HePO4, +Nisz (PO4) 2)

readily soluble in acids, even acetic acid.

12. Potassium Nitrite produces in dilute nickel solutions no pre-
cipitate (difference from cobalt). In very concentrated solutions a
brownish-red precipitate of Ni(NO)2-4KNOz is thrown down; in the
presence of alkaline earth salts a yellow crystalline precipitate is formed;
e.g., Ni(NO2)2:Ba(NO2)2:2KNOz2, which is very difficultly soluble
in cold water, but readily soluble in boiling water, with a green color.

18. Dimethylglyoxime. The reagent is prepared by dissolving 1 gm.
of the solid in 100 cc. of 98 per cent alcohol. If a little of the
reagent is added to a solution of a nickel salt, then ammonia to slightly
alkaline reaction, and the solution is boiled, a red crystalline precipitate
of the nickel salt of dimethylglyoxime is formed:

CH3—C = NOH
2

CH;—C=NOH

If the quantity of nickel present is very small, at first a yellowish
solution is obtained from which, on cooling, red needles are deposited.
According to L. Tschugaeff,* who first proposed this qualitative test,
the presence of one part of nickel can be detected in the presence of
400,000 parts of water. The reaction is not influenced by the presence
of ten times as much cobalt; when a larger proportion of cobalt is
present, the following procedure is followed,

4-NiClo-+-2NH3=2NH,Cl+ (CsH4N,0.) Ni.

Detection of Traces of Nickel in Cobalt Salts

Add strong ammonia to the solution of the cobalt salt until a clear solution is
obtained, then add a few cubic centimeters of hydrogen peroxide and boil the
solution a few minutes to decompose the excess of this reagent. Then add the
dimethylglyoxime and again bring the solution to a boil. A very small quan-
tity of nickel causes a red scum to form and the glass sides of the beaker become
coated with a film of red crystals. With smaller amounts of nickel the color
is best observed upon the filter through which the solution is poured and the
residue washed with hot water.

The above reaction is the most sensitive test known for detecting nickel
in the presence of cobalt.

v. Fortinit recommends the following method for detecting nickel in alloys:
Heat the metal, from which all greasy or oily matter has been removed, at one
place with the oxidizing flame of the blowpipe. Cool and moisten with one

* Ber. 38 (1905), 2520.
t Chem. Zitg.(1912), 1461.

178 REACTIONS OF THE METALS

drop of a solution of 0.5 gm. dimethylglvoxime in 5 cc. aleohol and 5 ec. concen-
trated ammonium hydroxide. A red spot at once appears if nickel is pres-
ent. When copper is present, the nickel test is obtained before the copper
ammonia color is visible.

REACTIONS IN THE DRY WAY
\

The borax, or sodium metaphosphate, bead is brown in the oxi-
dizing flame, almost the same shade as the strongly saturated man-
ganese bead; in the reducing flame the bead becomes gray, due to the
formation_of some metallic nickel. On looking at the bead through
the microscope the finely divided metal can be seen suspended in the
colorless glass. |

On heating nickel salts with sodium carbonate on charcoal a gray —

scale of metallic nickel is obtained. This reaction is best performed
with the charcoal stick, as described on p. 65. The magnetic metal
obtained in this way is placed on a piece of filter-paper, dissolved
in nitric acid, a drop of concentrated hydrochloric acid is added, and
the paper carefully dried by moving it back and forth over the flame.

If nickel is present the paper appears greenish (colorless with very small ©

amounts of nickel), or bluish if cobalt is also present. The paper is now
moistened (where the nickel is) with caustic soda or potash, and is then
held in bromine vapors, which are obtained by shaking some bromine
water in a wide-mouthed flask.

If nickel or cobalt is present, a black spot will be formed by the
above treatment, consisting of the hydroxide of the trivalent metal
(p. 174). The blackening often does not appear at first; in this case
the paper is moistened once more with potassium hydroxide and again
treated with bromine. The spot will now appear if nickel is present.

COBALT, Co. At. Wt. 58.97
Sp. Gr.=8.5. M. Pt. = 1490°
Occurrence.—Like nickel, native cobalt is found only in meteorites.

It occurs in the earth’s crust chiefly as sulfide, arsenide, and as salts
of thioarsenious and thioantimonous acids; it is almost always accom-

panied by nickel and iron. The most important ores are smaltite, —

CoAsz, isometric; cobaltite, CoAsS, isometric; skutterudite, CoAss,
isometric; erythrite, Cog(AsO4)2-8H20, monoclinic, isomorphous with
vivianite, Fe3(PO4)2-8H20, and with annabergite, Nis(AsO4)2:8H20.

Metallic cobalt is steel gray, dissolves much more readily in dilute
acids than nickel, and is, like the latter, magnetic. Cobalt forms, like
iron, three oxides: cobaltous oxide, CoO; cobaltous cobaltic oxide,
Co304; cobaltic oxide, Co2Qz.

ee is ee

COBALT 179

By dissolving these three oxides in acids, salts derived from cobalt-
ous oxide are always obtained, containing bivalent cobalt:

CoO +2HCl=H20 + CoCle,
Co203 +6HCl = 3H2O+ 2CoCle +- Cle t ;
Co304+8HCl=4H20+3CoCle+Cle fT .

Simple cobaltic salts are unknown, but many complex compounds
exist with trivalent cobalt, as, for example, potassium cobaltinitrite,
potassium cobalticyanide, and numerous cobalti-ammonia deriva-
tives.

- Cobaltous compounds in a crystallized state (as well as in aqueous
solution) are pink, in the anhydrous condition yellow or green, and
blue in aqueous solutions in the presence of hydrochloric acid. The
solubility reactions of cobaltous salts are similar to those of manganese
and nickel.

| REACTIONS IN THE WET WAY

1. Potassium or Sodium Hydroxide precipitates in the cold a blue

basic salt: |
Cott+Cl-+0OH- — Co(OH)CI,

which on warming is further decomposed by hydroxyl ions forming
pink cobaltous hydroxide: :

Co(OH)C1+OH- — Cl-+Co(OH)>.

In the case of a moderately concentrated solution of the alkali the
precipitate of pink cobaltous hydroxide is often produced in the cold,
sometimes only after standing some time. The rapidity of the reaction
depends entirely upon the concentration of the alkali.

Cobaltous hydroxide gradually turns brown in contact with the air,
going over into cobaltic hydroxide:

2Co(OH)2+H20+0 — 2Co(OH)s.

In this respect cobalt behaves similarly to iron and manganese,
and differs from nickel, for the hydroxide of the latter is not oxidized
by atmospheric oxygen.

On adding chlorine, bromine, hypochlorites, hydrogen peroxide,
etc., to an alkaline solution containing cobaltous hydroxide, cobal-
tic hydroxide is immediately formed, as with nickel and manganese,

2Co(OH)2+20H-+Cle = 2CI-+2Co(OH)s,
2CO(OH)2-+HOH+0Cl-=Cl-+2Co(OH)s.

180 REACTIONS OF THE METALS

From ammoniacal cobalt solutions the above oxidizing agents
cause no precipitation, but merely a red coloration; the addition of
potassium hydroxide then causes no precipitation (difference from
nickel).

Cobaltous hydroxide, Co(OH).2, behaves under some conditions as a weak
acid, for on adding to a cobaltous solution a very concentrated solution of KOH
or NaOH the precipitate at first produced dissolves with a blue color* similar
to copper. By the addition of Rochelle salts to this blue cobalt solution the
color either disappears almost entirely or becomes a pale pink, while the similarly
treated copper solution becomes more intensely blue. By the addition of
potassium cyanide to the blue cobalt solution it becomes yellow, and in contact
with air turns intensely brown. A copper solution would be decolorized by the
addition of potassium cyanide.

By pouring a little cobalt solution (or adding a little solid cobalt carbonate)
into a concentrated solution of caustic soda or potash, to which a little glycerol
has been added, a blue solution is formed (the color being intensified by warm-
ing), which after standing some time in the air, or immediately upon the addition
of hydrogen peroxide, becomes a beautiful green.

2. Ammonia precipitates, in the absence of ammonium salts, a
blue basic salt, soluble, however, in excess of ammonium chloride.
Ammonia, therefore, produces no precipitate in solutions which contain
sufficient ammonium chloride. The dirty-yellow, ammoniacal solution
is little by little turned reddish on exposure to the air, owing to the
formation of very stable cobalti-ammonia derivatives.

3. Alkali Carbonates produce a reddish precipitate of basic salt
of varying composition.

4. Ammonium Carbonate also precipitates a reddish basic salt,
soluble, however, in excess.

5. Barium Carbonate precipitates no cobalt in the cold and out
of contact with air, but on exposure to the air cobaltic hydroxide is
gradually thrown down. The precipitation takes place much more
quickly on the addition of hypocblorites or hydrogen peroxide:

2Cot ++2BaCO3+3H20+0CI — 2Ba*t+Cl-+ 2Co(OH)3+2CO2 fF

_ If the solution is heated to boiling, all of the cobalt is precipitated
as a basic salt, even out of contact with the air.

6. Hydrogen Sulfide produces no precipitate in solutions con-
taining mineral acids. In neutral solutions containing an alkali -
acetate all of the cobalt is precipitated as black sulfide.

7. Ammonium Sulfide precipitates black cobalt sulfide,

Cott+S8= — CoS,
* Ep. Donatn, Z. anal. Chem., 40 (1901), 137.

COBALT 181

insoluble in ammonium sulfide, acetic acid, and very dilute hydro-
chloric acid (cf. p. 175); soluble in concentrated nitric acid and aqua
regia, with separation of sulfur:

3CoS+8HNO3 =4H20+2NO 7 +3S+3Co(NOs)>o.

By continued action of strong nitric acid all the sulfur goes into
soiution as sulfuric acid.

The addition of an oxidizing agent always helps an acid to dissolve an insolu-
ble sulfide. The solution in contact with a sulfide precipitate at first contains
enough sulfide ions to satisfy the solubility product of the sulfide. When
hydrogen ions are added nonionized hydrogen sulfide is formed unless the
solubility-product of the sulfide is so small that less sulfide ions are present
than would be formed by the ionization of H.S. If an appreciable quantity
of HS is formed, it can be expelled as a gas and the sulfide will dissolve. Some-
times, however, this takes place very slowly and then the addition of an oxidiz-
ing agent is necessary. ‘The sulfide ions in solution are oxidized to free sulfur.
The solubility-product of the sulfide is no longer reached in solution, for as fast
as a little of the substance dissolves the sulfide ions are oxidized.

8. Potassium Cyanide produces in neutral solutions a reddish-
brown precipitate, soluble in excess of potassium cyanide in the cold,
with a brown color, forming potassium cobaltocyanide:

Cott+2CN~ — Co(CN)2; CoCN2+4CN- — [Co(CN)6¢]"==

On warming the brown solution for some time it becomes bright
yellow and reacts alkaline; it now contains potassium cobalticyanide,
of analogous composition to potassium ferricyanide. The formation
of the cobaltic salt takes place with the help of atmospheric oxygen:

2[Co(CN).6]~>~+O+H20 — 2[Co(CN).6]>+20H .

The reaction takes place more quickly in the presence of chlorine,
bromine, hypochlorites, etc.:

2[Co(CN)¢]*=+Cle — 2[Co(CN)¢]=+2CI-.

An excess of chlorine, bromine, etc., does not decompose the cobaltic
salt (difference from nickel).

The cobalticyanide anion is much more stable than the cobalto-
- eyanide anion. By adding hydrochloric acid to the brown solution of
potassium cobaltocyanide, hydrogen cyanide will be set free and yellow
cobaltous cyanide formed,

[Co(CN).]""+4Ht — 4HCN-+Co(CN)o,

while potassium cobalticyanide is not decomposed by hydrochloric acid.

182 REACTIONS OF THE METALS

Potassium cobalticyanide forms, with most of the heavy metals,
difficultly soluble or insoluble salts possessing characteristic colors.
Thus, it produces with cobaltous salts pink cobaltous cobalticyanide:

2[Co(CN)6]=+3Co** — Cos[Co(CN).le,

and with nickel salts greenish nickelous cobalticyanide.
If, therefore, a cobalt solution contains nickel it gives, on precipi- —

tating and redissolving with potassium cyanide, boiling, and adding

hydrochloric acid, a greenish precipitate of nickelous cobalticyanide:

2[Co(CN)6}" 4-3[Ni(CN)4] +12Ht — 12HCN+Nis3[Co(CN)glo.

9. Potassium Nitrite produces in concentrated solutions of cobalt
salts, with the addition of acetic acid, an immediate precipitation of
yellow, crystalline potassium cobaltinitrite. If the solution is dilute,
the precipitate appears only after standing for some time, but more
quickly on rubbing the sides of the beaker.

Cot t-+-7NO. +3Kt+2Ht — NO+H20 +Ks[Co(NO2) 6].

This reaction offers an excellent means for detecting the presence
of cobalt in nickel salts.

10. Ammonium Thiocyanate (Vogel’s reaction).* If a concen-
trated solution of ammonium thiocyanate is added to a cobaltous solu-
tion, the latter becomes a beautiful blue, owing to the formation of
ammonium cobaltothiocyanate:

Cott+4CNS~+2NH,4t > (NH4)2/Co(CNS)z].

On adding water the blue color disappears and the pink color of the
cobaltous salt takes its place. If amyl alcohol is added (or a mixture
of equal parts amyl alcohol and ether), and the solution shaken, the
upper alcoholic layer is colored blue. This reaction is so sensitive
that the blue color is recognizable when the solution contains only 74>
of a milligram of cobalt. The blue solution also shows a characteristic
absorption spectrum.{ Nickel salts produce no coloration of the amyl
alcohol. If, however, iron is present, red Fe(CNS)3 is formed, which
likewise colors the amyl alcohol, making the blue color (due to the

* Ber., 12, 2314; Treapwett, Z. anorg. Chem., 26 (1901), 105.

+ T. T. MorreEtt first showed that cobalt salts give a blue color with ammonium
thiocyanate, disappearing on the addition of water, but reappearing when alcohol
is added. Z. anal. Chem., 16, 251.

t Wotrr, Z. anal. Chem: 18, 58.

COBALT 183

cobalt) very indistinct, so that, under some conditions, it can no longer
be detected. If, however, 2 or 3 cc. of concentrated ammonium acetate
solution and 2 or 3 drops of 50 per cent. tartaric acid solution are added,
the red color produced by Fe(CNS)3 will disappear and the blue color
of the cobalt compound will be seen.

The blue color is probably that of undissociated (NH,)s[Co(CNS),]. When
the solution is diluted, the salt is ionized and the complex anion also is in
equilibrium with cobalt ions, the more dilute the solution, the greater the
ionization (cf. p. 19). The alcohol and ether probably dissolve only the un-
dissociated (NH,).[Co(CNS),] and this is evidently present to some extent in
the aqueous solution, although its color is obscured by that of cobaltous ions.

11. Ether Saturated with Hydrogen Chloride does not precipitate
an anhydrous cobaltous salt, as in the case of nickel, but will dissolve
‘the blue, anhydrous cobaltous chloride. This furnishes a basis of a
method for separating nickel and cobalt.

12. a-Nitroso-6-naphthol, CjoHe6(NO)OH, produces a voluminous,
purple red precipitate of cobalti-nitroso-8-naphtol, [Ci9H¢(NO)O]3Co,
which is insoluble in cold, dilute nitric or hydrochloric acid.*

This reagent serves not only for qualitative purposes, but can also be used
for the quantitative determination of cobalt in the presence of nickel: The
test may be applied conveniently to the solution obtained in the usual qualita-
tive scheme after the removal of all metals except nickel and cobalt.. A part
of the solution may be used for the sensitive nickel test with dimethylglyoxime
(p. 177) and the remainder used for the cobalt test.

Dilute the solution to about 50 cc., add 4 cc. of 6N hydrochloric acid and
20 ec. of 6N acetic acid. Heat, add 50.cc. of a saturated solution of nitroso-
6-naphthol and boil in 50 per cent acetic acid. If as much as 0.1 mg. of cobalt
is present, a red precipitate or turbidity is obtained even in the presence of 250
mg. of nickel. When more than 150 mg, of nickel is present, however, some of
the brownish-yellow nickel compound, (CioHs(NO)O).Ni, will precipitate after
the solution cools.

The reagent used in this test should be freshly prepared. Nitroso-
p-naphthol gradually decomposes on standing in the air and changes from yellow
to brown or even black in color. It can be purified by dissolving in hot sodium
carbonate, filtering and reprecipitating with sulfuric acid. For all ordinary
purposes the saturated solution in 50 per cent acetic acid is most suitable.
The cobalt test can be made more delicate by adding an equal volume of alcohol
to the test and, for detecting traces of cobalt, an aqueous solution of the organic
substance can be used, but as 5000 cc. of water are required to dissolve 1 gm. of
the nitroso-s-naphthol, it is evident that the aqueous solution is not suitable
when much cobalt is present. An excess of the reagent is required, as a part of
it is used to oxidize the cobalt to the trivalent condition.

Copper gives a characteristic coffee-brown precipitate with the reagent
and it is possible to separate copper from leagl, cadmium, etc., by means of it.

*TLInskI and v. Knorre, Ber., 18, 699 (1885),

184 REACTIONS OF THE METALS

Ferric iron gives a brownish-black precipitate which is insoluble enough to
serve as a means of separating iron from aluminium, manganese, etc. Ferrous
iron also gives a greenish precipitate in neutral solutions. Of all these precipi-
tates, however, the cobalt compound is the most characteristic and the least
influenced by the presence of acid. Thus with the acidity recommended above,
the presence of a little ferric or ferrous iron causes no disturbance.

Detection of Traces of Cobalt in Nickel Salts

To test a nickel salt for ‘cobalt, add a concentrated solution of
ammonium thiocyanate to the solution of a considerable amount of
the salt, a few cubic centimeters of a mixture of amyl alcohol and ether
and shake the mixture. After the latter has been allowed to settle,
if the upper alcohol-ether layer is colorless, then the nickel salt contains
neither iron nor cobalt; if the layer is reddish, iron is present. In the
latter case add 2 or 3 cc. of concentrated ammonium acetate solution
and 2 or 3 drops of 50 per cent tartaric acid solution and shake again;
if cobalt is present the alcohol-ether layer is now distinctly blue.

Sometimes when very little cobalt and considerable nickel is present it is
hard to tell whether the amyl alcohol is colored blue or not. In such a case
pour the solution into a separatory funnel and draw off the lower layer contain-
ing the green nickelous solution. Add a little more ammonium thiocyanate
solution to the amyl alcohol, 1 cc. of ammonium acetate solution, 1 drop of

tartaric acid solution and shake again. The blue color should now appear if
any cobalt is present.

REACTIONS IN THE DRY WAY

The bead produced by borax or sodium metaphosphate is blue in
both the oxidizing and reducing flames. By holding the bead in the
upper reducing flame for a long time it is possible to reduce the cobalt
to metal, when it appears, like nickel, gray.

On the charcoal stick cobalt compounds yield gray metallic cobalt,
which can be removed by means of a magnetized knife-blade, as
described on p. 66, placed on filter-paper, dissolved in hydrochloric
acid and dried. ‘The paper is then colored blue by cobalt (difference
from nickel). If, now, sodium hydroxide is added and the paper
exposed to the action of bromine vapors, black cobaltic hydroxide,
Co(OH)s3, is formed.

ZINC, Zn. At. Wt. 65.37

Sp. Gr.=6.9. M. Pt.=419°. B. Pt.=916°.

Occurrence.—Smithsonite, ZnCOg, isomorphous with calcite,
CaCOs, etc.; sphalerite, ZnS, isometric; calamine, Zn2Si04+He20,
orthorhombic, hemimorphic; zincite, ZnO, hexagonal; and franklinite,
(FeO2)2(Fe,Mn,Zn), isometric. :

ZINC 185

The most important zine ore is sphalerite, ZnS. Sulfide of zine
is dimorphic and is also found as wurtzite, which crystallizes in the
hexagonal system.

Metallic zine is bluish white. At low temperatures and at about
200° C. it is so brittle that it can be pulverized, but at 110°-150° C.
it is ductile and can be drawn out into wire and rolled into foil.

Zinc, as its position in the electromotive series would indicate (p. 41)
dissolves readily in all acids; in hydrochloric, sulfuric, and acetic
acids with evolution of hydrogen:

Zn+2H* > Zn**+Hp f.

Zinc is such a strong reducing agent that it easily reduces nitric acid,
the extent of the reduction depending upon the concentration of the
acid. With very concentrated acid, some NOz is obtained, while dilute
acid is reduced to ammonium nitrate. In concentrated acid the prin-
' cipal product is nitric oxide, NO.

8Zn+8HNO3 — 3Zn* *+6NO3° +4H20+2NO f¢.
4Zn+10HNO3 — 4Zn* *+8N03°+NH4NO3+3H20.

- Like aluminium, zine dissolves in caustic soda or potash, with
evolution of hydrogen and the formation of a zincate:

Zn+20H~— ZnOo-+He 1T.

Zine forms only one oxide, ZnO. It is a white infusible powder,
which becomes yellow when heated, but turns white again on cooling.
Zinc oxide dissolves readily in acids, forming zinc salts:

ZnO +H22SO4 = H20 +7ZnSOx.

There exists only one series of zinc salts, and the zinc is always
bivalent. Most of the salts are white. The chloride, nitrate, sul-
fate, and acetate are soluble in water; the remainder dissolve readily
in mineral acids.

REACTIONS IN THE WET WAY

1. Potassium or Sodium Hydroxide precipitates white, gelatinous
zinc hydroxide, easily soluble in excess of the precipitant, forming a
— gzineate:*

Zn+*++20OH- — Zn(OH)2; Zn(OH)2+OH- @ HZnO2 +H20.

* According to Hantzscu the zinc is not present as zincate, but probably in
colloidal solution. Z. anorg. Chem., 30, 289 (1902). In fairly concentrated solu-
tions, however, it is certain that the zinc is present as zincate, for F. Forrster
and O. Ginruer, Z. Electrochem., 6, 301 (1900), have isolated the compound,
NaHZnO -3H,0, as needles with silky luster.

186 REACTIONS OF THE METALS

Zinc hydroxide, therefore, behaves sometimes as a base and some-
times as an acid, like aluminium hydroxide.

On boiling a diluted solution of a zincate, hydrolysis takes nlids
and zine hydroxide is precipitated, but if the solution contains an excess
of OH ions, there will be no precipitation.

2. Ammonia precipitates from neutral solutions, free ‘from am-
monium salts, zinc hydroxide, readily soluble in ammonium galts,
as in the case of magnesium, nickel, manganese, or iron:

Zn* ++9NHsOH ra 4 Zn(OH)>2 +2NHa,".

Zinc hydroxide is also soluble in an excess of ammonia, due to the
formation of complex zinc ammonia ions:

Zn(OH)2+6NH3 — [Zn(NH3)6]* t+20H-.

3. Alkali Carbonates precipitate a white, basic carbonate, of
variable composition, as is the case with magnesium.

4. Ammonium Carbonate does the same, except that the pre-
cipitate is soluble in an excess of the reagent. ‘The presence of am-
monium salts or of ammonia prevents the precipitation.

5. Barium Carbonate precipitates no zinc in the cold, but on boiling
all the zinc is precipitated as basic carbonate.

6. Sodium Phosphate precipitates gelatinous, tertiary zine phos-
phate, which soon becomes crystalline, and is soluble in ammonia and
in acids:

Zn* *++4HPO« — ZHePO4- +Zn3(PO4)o.

In the presence of ammonium salt, the less soluble zinc ammonium
phosphate is precipitated:

Zn* +4+4.NH,4t+2HPO," — HePO, +ZnNHy4POs,.

Both zine phosphate and zinc ammonium phosphate dissolve readily
in dilute acids, owing to the formation of very slightly ionized HPO4’,
and in dilute ammonia, owing to the formation of zinc ammonia cations. °
Acids, therefore, deprive the solution of PO4> anions by forming HPO’,
and ammonia deprives the solution of zincions by forming [Zn(NH3)¢6]**

7. Hydrogen Sulfide precipitates the zine as sulfide, from neutral
solutions of a zinc salt:

Zn+++H.S = ZnS+2H".

The solubility product of zinc sulfide (p. 22) is about 1.2X107%%. At 25°
the concentration of a saturated solution of hydrogen sulfide is about 0.1 molar

ZINC 187

and the ionization constant for the complete ionization, H.S = 2H-+5S, has been
estimated to be 1.110~*8. The concentration of the sulfide ion in such a solu-
tion is approximately 1.2X10~ molar equivalents per liter. The solubility-
product of zine sulfide is evidently exceeded when the aqueous solution of
a zinc salt is saturated with hydrogen sulfide and zine sulfide is precipitated.
The mass-action principle applied to the complete ionization of hydrogen
sulfide shows that the concentration of sulfur ions is inversely proportional
- to the square of the concentration of the hydrogen ions. If the concentration
of the hydrogen ions is increased one thousandfold, and this is approximately the
case when the solution is tenth-normal with a mineral acid, the concentration
of the sulfide ions from hydrogen sulfide is reduced to one-millionth of its value
in pure water.

The separation of the second group of metals from the third group is usually
accomplished by passing hydrogen sulfide into a solution which is about 0.3-
normal with hydrochloric or nitric acid. The concentration of the sulfide ion
when such a solution is saturated with hydrogen sulfide at 25° is about
1.1X10-*. To reach the solubility product of zine sulfide in 0.3-normal acid,
the zinc ions should reach the concentration of about 0.11 mole per liter, or
about 0.7 gm. per 100 ce. .

If the zinc salt has a greater concentration than this, some zine sulfide
should be precipitated by hydrogen sulfide in 0.3-normal acid solution. The
precipitation would evidently be incomplete and, as more hydrogen ions are
formed in solution from the hydrogen sulfide as a result of the sulfide precipita-
tion, the ionization of the hydrogen sulfide continually tends to become less.

If, however, considerable sodium or ammonium acetate is added to the acid
solution, the concentration of the hydrogen ion becomes much smaller and the
ionization of the hydrogen sulfide takes place to a greater extent. It is then
possible to precipitate the zinc as sulfide so completely that less than one i mg.
of zinc will remain in solution.

Zinc sulfide dissolves readily in normal hydrochloric acid. The sulfur ions
from the zinc sulfide enter into equilibrium with the hydrogen ions of the acid
to form hydrogen sulfide. In normal acid solution, the concentration of sulfur
ions from saturated hydrogen sulfide is about 1.2X10~%* and of sulfur ions
from a saturated solution of ZnS in water about 3.5107".

8. Ammonium Sulfide precipitates from neutral or alkaline solutions
all the zine as amorphous sulfide:

Zntt++S"- —> ZnS.

Zinc sulfide is a precipitate hard to filter; it runs through the
filter-paper, particularly on washing. This is a peculiarity of almost
all metallic sulfides and of many other amorphous bodies, such as |
aluminium hydroxide, titanic acid, tungstic acid, and many others.
This is due to its tendency to form colloidal solutions (p. 58). The
colloid can be precipitated by adding a concentrated salt solution or
by boiling.

* This value is merely an approximation, being derived by a rough calculation
from values which are not very reliable.

188 REACTIONS OF THE METALS

In order, then, to obtain zine sulfide in a form which can be
filtered, it is best precipitated from a boiling solution containing acetic
acid and a considerable quantity of ammonium salts. The precipitate
may be washed with a solution of ammonium chloride to which a little
ammonium sulfide has been added.

9. Potassium Cyanide produces a white precipitate of zine cyanide,
soluble in an excess of the precipitant.

Znt*++2CN~ @ Zn(CN)2; Zn(CN)2+2CN7~ — [Zn(CN)a]*.
The zine-cyanide anion is decomposed by acids and by alkali sulfide:
[Zn(CN)4]-+2Ht — Zn(CN)2+2HCN;
[Zn(CN)4]- +S" — ZnS+4CN-.

10. Potassium Ferrocyanide precipitates white zinc ferrocyanide,
which is changed by-an excess of the potassium ferrocyanide into less
soluble zinc-potassium ferrocyanide:

¢ [Fe(CN) 6]=~+2Zn* + — Zno[Fe(CN)6]
8Zn2[Fe(CN)«]+Ks[Fe(CN)¢] — 2K2Zn3[Fe(CN)6]2

REACTIONS IN THE DRY WAY

Heated with sodium carbonate on charcoal before the blowpipe,
it is not possible to obtain metallic zine on account of its volatility;
but an incrustation of oxide is obtained which is yellow while hot
and white when cold.

Zine oxide (or such compounds of zine as are changed over to oxide
on ignition), when moistened with cobalt nitrate yields a green infusible
mass—Rinnmann’s green. This reaction is performed exactly as with
aluminium (p. 131).

Separation of the Metals of Group III from the Alkalies and
Alkaline Earths

The separation of the members of the ammonium-sulfide group
from the alkalies and alkaline earths is effected by means of ammo-
nium sulfide and ammonium chloride. If, however, the solution
contains phosphoric acid, oxalic acid, or considerable boric acid, the
neutralization of the solutions will cause the precipitation of calcium,
strontium, barium, and magnesium as phosphate, oxalate, or borate.
The procedure to be followed when such acids are present will be given
in Part IV of this book after the characteristic properties of the igi
have been described in Part III.

Taste V.—ANAtysis or Group III 1n ABSENCE OF PHOSPHATE.

ANALYSIS OF GROUP III

189

Metuop A

Solution may contain: Fett+, Fet+,UO,++, Altt+, Crt++, Mnt+, Znt+,
Cot*+, Nit*+, also Groups IV and V. Add NH,OH and (NH;,)S. Filter and

examine filtrate for Groups IV and V. Treat ppt. with 2-normal HCl. (2)
Residue: CoS, | Solution: Fet+, UO.++, Al+++, Crt++, Mn++, Zn++ [and some
NiS. Test for | Ni]. Add NaOH in excess. (4)
Ni and Co by
bead test. Test | precipitate: Fe(OH);, Cr(OH)s, | Filtrate: AlO.-, HZnO.-. Add
for Ni with) Na,U.07, Mn(OH)2 [and some| HCl and then NH,OH. (10)
Nie de an Ni(OH).]. Dissolve in HCl,
find for cobalt add NH,Cland NH,OH. (5)
eae Roar Precipitate:|Filtrate: |Preci pi -| Filtrate: Zntt.
anate. (3) Fe(OH);, Cr(OH);,}, Mn++] tate: AI(OH);| Acidify with
(NH4)2U207. Dis-| (and Confirm by| HC.H;02. add
solve in HCl| Nitt)|] Thénard’s| HS and con-
and add excess| Test for| bluetest. (10)| firm Zn _ by
(NH4)2CO3. (6) Mn by Rinnmann’s
(NH4)2S green test.
Precipitate:|Filtrate: et de (11)
Fe(OH )s, | (NHs)U02(COs)s} Gp py fun
Cr(OH)s.| Acidify with) bith
Test for| HCl and test NaCO
Fe with| with SS.
K.iFe (CN )e. K4Fe(CN )<¢. KNO
TestforCrby | (8) (9) ;
fusion with
Na,CO; and
KNOs, ete. (7)

A number of excellent schemes have been devised for the analysis of this
group and each has something in its favor. In this book it has seemed best
not to attempt to decide upon any one scheme, but rather to treat the subject
in a broader manner, partly because of the instructive value of studying several
schemes and partly because one scheme is best under certain conditions and
another scheme under different conditions. Thus it is a quite common practice
to divide the whole ‘group into two minor groups, one containing ferric iron,
aluminium, chromium, and uranium and the other containing manganese,
nickel, cobalt and zinc. Such a scheme often works very nicely and enables one
to arrive at proper conclusions quickly but, unfortunately, chromium when
precipitated in this scheme has a marked tendency to carry down zine and
magnesium with it, and this may result in the failure to detect zinc or magnesium
during the subsequent examination. Chromium, however, on account of the
color of its compounds, invariably betrays its presence before the actual test
for chromium is made, and it is a very easy matter to modify the method some-
what when chromium is present in order that zinc and magnesium will not be
missed. It is unnecessary to use such a modified method when chromium is
known to be absent. In most schemes of analysis the detection of uranium and
titanium is not provided for in the analysis of this group. Titanium, although

190 REACTIONS OF THE METALS

present in most rocks, is usually found in very small quantities. When much
titanium is present, this fact is known by the difficulty involved in getting the
substance in solution and the tendency of the dilute acid solutions to hydrolyze
and form precipitates of metatitanic acid, which are difficult to filter. When
much titanium is present it is precipitated according to p. 158, §4. To detect
small quantities of titanium, the hydrogen peroxide test (p. 159) is most suit-
able. Uranium is of relatively rare occurrence, but its ores have become
important since the discovery of radium in them. The detection of uranium
will be included in the first of the schemes of analysis that follow.

PROCEDURE

1. Heat the neutral solution to boiling, add 5 ec. of normal ammonium
chloride solution, if this salt is not already present, and ammonium sulfide
solution drop by drop until no further precipitation takes place. Avoid add-
ing an excess of ammonium sulfide on account of the danger of getting a turbid
precipitate when nickel is present. (To avoid this danger it is well to pass
hydrogen sulfide into the slightly ammoniacal solution instead of adding ammo-
nium sulfide. The reason why nickel sulfide runs through the filter is partly .
because ammonium polysulfide is present to some extent in the ammonium sul-
fide reagent that is not freshly prepared.) Filter off the precipitated sulfides
and wash promptly with hot water. If the moist sulfides are allowed to stand
exposed to the air, some sulfate is formed by oxidation and this will dissolve in
the wash water. Reject all but the first washings and use the filtrate for the
analysis of the alkaline earth and alkali groups (p. 111).

2. Digest the precipitated sulfides in a porcelain dish with cold, 2-normal
hydrochloric acid, stirring until no more hydrogen sulfide is evolved. Filter
off the residue, which consists chiefly of cobalt and nickel sulfides and wash with
a little hydrochloric acid. Usually a partial oxidation of the sulfide takes place
during this treatment with dilute acid and sulfur is formed which is likely to
enclose a little sulfide that should dissolve in the acid. The fact that a residue
remains is not, therefore positive proof of the presence of nickel or cobalt.
Examine the solution by § 4.

3. Test the residue for cobalt by heating a little of it in a borax bead; a
blue bead shows cobalt. If the borax bead is brown, further test for nickel is
unnecessary. If a blue bead was obtained, test for nickel by dissolving the
precipitate in aqua regia, evaporating just to dryness, adding a little water
and testing with dimethylglyoxime (p. 176). If a brown bead was obtained,
dissolve the precipitate as just described and, to the solution freed from mineral
acid, add a concentrated solution of potassium nitrite, acidify with acetic acid
and allow the solution to stand at least ten minutes. A fine yellow precipitate
of K;[Co(NO:).] shows that cobalt is present.

4, Evaporate the solution obtained in (2) to a small volume, oxidize any
iron present to the ferric condition by heating with a little strong nitric acid,
then add sodium hydroxide solution until a strongly alkaline solution is obtained,
boil and filter. Examine the filtrate by § 10. ‘«

The precipitate may contain iron, chromium, uranium and manganese
(with a little nickel) and the filtrate may contain aluminium and zine.

5. Dissolve the precipitate in as little hydrochloric acid as possible, dilute
the solution with hot water and boil several minutes. Add 5 cc. of ammonium

ANALYSIS OF GROUP III 191

chloride solution, make barely alkaline with ammonia and filter promptly. The
precipitate contains all the iron, chromium and uranium; the filtrate, which
may contain manganese, and traces of nickel, is analyzed by § 9.

6. Dissolve the precipitate in as little hydrochloric acid as possible, add a
large excess of ammonium carbonate solution, heat gently but do not boil long,
and filter. The precipitate contains the iron and chromium; the filtrate
contains uranium in solution as ammonium. uranyl carbonate. Analyze
by § 8.

7. Test the precipitate obtained in (6) for iron by dissolving a part of it in
a few drops of hydrochloric acid, diluting with a little water and adding potas-
sium ferrocyanide solution. The formation of Prussian blue shows the presence
of iron. Test another portion of the precipitate for chromium by mixing it
with sodium carbonate and potassium nitrate and fusing to form sodium
chromate. Dissolve the melt, which is yellow if chromium is present, in water,
acidify the aqueous extract with acetic acid and add a drop of silver nitrate
solution; a red precipitate of silver chromate is formed if chromium is present.

8. To test for uranium, add hydrochloric acid to the solution obtained in (6)
and treat the slightly acid solution with potassium ferrocyanide; a brown pre-
cipitate shows the presence of uranium.

9. Test the filtrate from (5) for manganese. Evaporate the solution to
dryness, dissolve the residue in a little water and add a few drops of potassium
cyanide solution. Dilute with water, add ammonium sulfide and boil. A
flesh-colored precipitate is MnS. To confirm the test, dissolve a portion of the
precipitate in concentrated nitric acid, add a little lead peroxide and boil.
Dilute with water and allow the precipitate to settle; manganese is shown by
the characteristic color of the permanganate ion. Or, a small portion of the
manganous sulfide can be fused with sodium carbonate and potassium nitrate.
A green melt shown shows the presence of manganese.

10. Test the filtrate obtained from (4) for aluminium and zinc. Make it
acid with hydrochloric acid and then add a slight excess of ammonia. A white
precipitate is Al(OH);. Filter and test the filtrate for zinc (11). To confirm
the aluminium test, dissolve the precipitate or a small portion of it in nitric acid
and add half as many drops of 1 per cent cobaltous nitrate solution as there
are presumable milligrams of aluminium in the precipitate and evaporate the
- solution nearly to dryness in a casserole. Soak up the solution in a small piece
of filter paper, roll up the paper and wind a platinum wire around it. Heat in
the flame till all the paper is consumed and then ignite strongly. Thénard’s
blue shows the presence of aluminium (p. 1381).

11. To test for zinc, acidify the solution obtained in (10) with acetic acid and
saturate it with hydrogen sulfide. Filter and dissolve the precipitate in a little
nitric acid. Add 1 drop of é¢obalt nitrate solution and as many more drops as
there are estimated to be centigrams of zinc present. Evaporate to dryness
and ignite the contents of the dish until the purple color of the cobalt salt disap-
pears, Rinnmann’s green (p, 188) shows the presence of zinc,

192

REACTIONS OF THE METALS

TaBLE VI.—ANAtysis or Group III in ABSENCE OF PHOSPHATE
Method B

Solution may contain Fe++, Fet++, Al+++, Crt++, Mn++, Zntt, Cott, Nit+,
and Groups IV and V (p. 111 or p. 118).
Add NH,OH and (NH,)S._ Filter and examine filtrate for Groups IV and V.

Dissolve ppt. in HCl and HNO3.

Evaporate, treat with NaOH and Na,O2 and —

filter. (2).

Precipitate: Fe(OH);, H2MnO;, Co(OH)s,|| Filtrate: AlO.-, CrOs,
Ni(OH).[Zn(OH),]. Dissolve in HNO; and H.2O2.|| HZnO.-. Acidify with HNO;
Evaporate and boil with conc. HNO3;+KCI103. (8) || and add NH,OH,. (12)

Precipi-|Filtrate: Fet++, Cot+, Nitt, [Znt+]. ||/Precipi-|Filtrate: CrO,, Zn.

tate:| AddNH,OH. (5) tate :| Add HC.H;0, and

MnOz AIOH);.| BaClh. (14)

LEpeoesa Precipi- [Filtrate: Co(NHs)st+t, Ape

+tate: Ni(NHs).tt, [Zn(NHs3)¢6t+ +]. al .
+ 202) Pe(OH);.| Saturate with HyS and treat ppt. metho dt Tecipi- ete
and test . A. (13)} tate: Zn. Sat
Test for| with 2-normal HCl. (7) BaCrO,: |/sanaee

for Mn Fe as in Di 3 ith H.

with | Method|pecdue: ; + eae |

NaBiOs.| ‘4 (G Residue: |Solution: Zntt, in H Cl and con-

(4) - (6) CoS,NiS.} traces of Cott and and | firm zine

Test for| Nit+. Add NaOH H,.SO3.| asin
Ni and} and Na,O2. (9) Evapo- | Method
Co as in rate. A. (16)
Method |py eeipi-|Filtrate: Green
A. (8) | tat e:| NasZnOp. color
Co(OH):| Acidify shows
; Ni(OH)s.| with Cr. (15)
Add to| HC,H;0,
residue | and sat-
of COS,) urate
NiS.(10)| with HS
Confirm:
Zn as in
Method
A. (11)
PROCEDURE :

1. Precipitate with ammonium sulfide as in Method A, filter and examine

the filtrate for the alkali and alkaline-earth metals.

2. Digest the sulfide precipitate with hot, 6-normal hydrochloric acid anit

add enough nitric acid to dissolve the nickel and cobalt sulfides.

water and filter off the residual sulfur.
ness to remove the excess of acid, dilute to about 25 ce. and carefully neutralize
If a very heavy precipitate is produced, it is
To the cold solution carefully add a

with pure sodium hydroxide.
best to dilute with a little more water.
little sodium peroxide powder.

ilute with

Evaporate the solution nearly to dry-

(On account of the violent reaction with water,

and the fact that the powder often contains a little free sodium, care should be

ANALYSIS OF GROUP III 193

taken not to add the peroxide too fast or to a hot solution. Only a little perox-
ide should be taken from the container at one time and it should be transferred
directly to glass and never to paper.) Finally boil the solution to decompose
the excess of peroxide, dilute with an equal volume of water and filter. The
precipitate contains ferric hydroxide, hydrated manganese dioxide, cobaltic hydrox-
ide and nickelous hydroxide. The filtrate contains sodium aluminate, chromate
and zincate. ‘The separation is faulty in the case of zine which normally stays
in solution. “As much as 5 mg. of zinc may be carried down with the precipitate
when much iron, nickel or cobalt is present, and as much as 20 mgs. by con-
siderable manganese. This is probably due to the amphoteric nature of the
precipitated hydroxides and the insolubility of the zinc salts of the corresponding
acids. Examine the filtrate for chromium, aluminium and zinc by § 12.

3. Dissolve the precipitate in hot, 6-normal nitric acid, adding as much hydro-
gen peroxide as is necessary to reduce the manganese and cobalt to the bivalent
condition (cf. p. 33). Evaporate the solution nearly to dryness, add 15 ce.
of 16-normal nitric acid and about 1 gm. powdered potassium chlorate and heat
_ to boiling. Add 10 cc. more of concentrated nitric acid, heat to boiling, remove
the flame and add 0.5 gm. more of potassium chlorate. Repeat the treatment
with fresh portions of chlorate until about 3 gms. of chlorate have been used. Do
not add the chlorate to the nitric acid solution while it is boiling, as an explosion
is likely to result, but boil after the addition of each portion of chlorate. The
treatment with chlorate is best accomplished in a 250-cc. Erlenmeyer flask.
If a precipitate of MnO, is formed, filter through a thin layer of good-quality
washed asbestos which is supported by a little glass wool in an ordinary funnel.
Test the filtrate for manganese by adding 1 gm. of potassium chlorate and boil-
ing again. Wash the precipitate with a little concentrated nitric acid, which
has been freed from nitrous acid by boiling with a little potassium chlorate
just previous to use. Examine the filtrate for iron, cobalt, nickel and zine
by § 5.

4. Dissolve the precipitated manganese dioxide with a little hot 6-normal
nitric acid and a few drops of hydrogen peroxide. Boil to decompose any excess
of the latter and cool to room temperature. Add a little solid sodium bismuth-
ate, shake and let the solid settle. A purple solution shows the presence of
manganese (cf. p. 167).

5. Add an excess of ammonia to the filtrate from (3) to precipitate ferric
hydroxide, leaving cobalt, nickel and possibly some zinc in solution as soluble
complex metal-ammonia cations. Examine the filtrate by § 7.

6. Examine the precipitate for zron as in Method A.

; 7. Saturate the filtrate from (5) with hydrogen sulfide, filter off the precipi-

tated cobalt, nickel and zine sulfides and reject the filtrate. Digest the pre-
cipitated sulfides with 2-normal hydrochloric acid to dissolve any zine sulfide
that may be present; a little nickel and a trace of cobalt may be dissolved by’
this treatment. Examine the solution for zine by § 9.

8. Test the sulfide residue for nickel and cobalt as in Method A.

9. Neutralize the solution obtained in (7) with sodium hydroxide and add
sodium peroxide as in (2). Examine the filtrate by § 11.

10. If deemed advisable, any precipitated Ni(OH). and Co(OH); may be
added to the sulfide residue obtained in (7) and tested for nickel and cobalt as
in (8).

11. Acidify the filtrate from (9) with acetic acid and saturate the solution

194 REACTIONS OF THE METALS

with hydrogen sulfide. Any precipitate that forms is probably zine sulfide.
Confirm the zinc test as in Method A.

12. Acidify the filtrate from (3) with nitric acid and add ammonia until
present in slight excess. Heat to boiling to coagulate any precipitated alumin-
ium hydroxide and filter. Test the filtrate by § 14.

13. Confirm the presence of aluminium as in Method A.

14. If chromium is present in the filtrate from (12) it is shown by the
yellow color of the chromate ions. If the solution is colorless at this point it
is unnecessary to test for chromium. If it is yellow, carefully neutralize with
acetic acid until a slight excess is present, heat to boiling and precipitate by
the gradual addition of hot barium chloride solution. Filter and test the
filtrate for zinc by § 16. ;

15. The yellow precipitate of barium chromate is conclusive evidence of
the presence of chromium. Sometimes the yellow precipitate is obscured by a
white precipitate of barium sulfate. To confirm the chromium test, dissolve the
precipitate of barium chromate by pouring a mixture of 3 cc. 6-normal hydro-
chloric acid and 10 ec. of saturated sulfur dioxide solution through the filter
several times. LEvaporate the filtrate to dryness in a porcelain dish, taking
care not to overheat the residue, and add a few drops of water. A green color
shows the presence of chromic ions which were formed by the reduction of the
chromate. Sometimes a yellow color is obtained during the evaporation. This
is due to the presence of a very little ferric chloride which has gotten into the
solution accidentally. It does not interfere seriously with the test.

16. Saturate the filtrate from (14) with hydrogen sulfide and if a pe
of zine sulfide is formed confirm the test as in Method A.

METALS OF GROUP II. HYDROGEN SULFIDE
GROUP

MERCURY, LEAD, COPPER, BISMUTH, CADMIUM, ARSENIC, ANTIMONY,
TIN (GOLD, PLATINUM’

MERCURY, Hg. At. Wt. 200.6
Sp. Gr.=13.60. M. Pt.=—38.7°. B. Pt.=357°

Occurrence-—Mercury occurs in nature chiefly in the form of
rhombohedral cinnabar, HgS; from the ore, free mercury is obtained
by sublimation. According to G. F. Becker,* cinnabar is deposited
from solutions of its thio salt. The richest deposits are those of New
Almaden in California, where it occurs with serpentine; of Almaden
in Spain, Idria in Carniola, and Moschellandsberg in the Palatinate of
the Rhine. With cinnabar small quantities of native mercury are
often found. Mercury is also an important constituent of many varie-
ties of tetrahedrite.

Metallic mercury is the only one of the metals which is liquid at
ordinary temperatures. It is insoluble in hydrochloric and dilute sul-
furic acids, but is soluble in hot concentrated sulfuric acid with evolu-
tion of sulfur dioxide, forming mercurous or mercuric sulfate accord-
ing to whether the metal or the acid is present in excess:

Hg+2H2804 = HgSO4+2H20+S80z2 fT
QH¢+2H2SO04 = Hge8O1+2H20+802 ft

Hydrobromie acid hardly attacks the metal at all, while in hydriodic
acid the metal dissolves readily with evolution of hydrogen:

Hg+4HI = H2[Hgl4]+He Tt

The position of mercury in the electromotive series (p. 41) shows
that mercury cannot be oxidized by H* except when the concentration
of Hg**t is extremely low. This explains why mercury does not dissolve
in dilute hydrochloric or sulfuric acid. It seems remarkable, therefore,
that mercury should be oxidized by hydriodic acid. The reason the

* Geology of the Quicksilver Deposits of the Pacific Slope. Washington, 1888.
195

196 HYDROGEN SULFIDE GROUP

hydrogen of hydriodic acid can accomplish the oxidation of the mercury
is because the compound H2{HglI4] is scarcely dissociated at all into
Hg** ions (ef. p. 10).

The proper solvent for mercury is nitric acid.

If the metal is treated with hot concentrated nitric acid, mercuric
nitrate is formed:

3Hg+8HNO3=3Hg(NO3)2+4H20+2NO fT .

If, however, cold nitric acid is allowed to act upon an excess of mer-:
cury in the cold, mercurous nitrate is obtained:

Hg(NOs) 2 +Hg = Hge (NOs) 2.

Mercury is attacked by chlorine, forming calomel (mercurous
chloride) :
2H¢g+Cle = HgeCle.

Two oxides of mercury are known: yellow or red mercuric oxide,
HgO; black mercurous oxide, Hg2O. |

These oxides are basic anhydrides, from which two series of salts
are derived: (a) The mercuric salts, which contain Hgt*, and (b) the’
mercurous salts, which contain the group Hget*. We will consider
first the more stable mercuric salts.

Mercuric Salts

Mercuric salts are mostly colorless. The iodide is red or yellow.
By heating the red tetragonal crystals of mercuric iodide a yellow sub-
limate of orthorhombic needles is obtained, which gradually changes
back to the red tetragonal modification; very quickly, almost instantly
if the yellow crystals are rubbed. This is a general property of dimor-
phous bodies; the more symmetrical form is almost always the more
stable.

The sulfide is black or red.

Mercurie chloride is soluble in water, 100 cc. of water dissolving
6.57 gm. at 10°, 7.39 gm. at 20 cc., 11.34 gm. at 50 cc., 24.8 gm. at
80° and 53.96 gm. at 100°.

In water containing hydrochloric acid, mercuric chloride is much
more soluble than in pure water; and in fact the solubility increases
with the concentration of the hydrochloric acid, due to the formation
of the complex acid H2[HgCl4]. Alkali chlorides also help ‘to dissolve
mercuric chloride, forming salts of this complex acid. Mercuric chloride
is more soluble in alcohol and in ether than it is in water.

The aqueous solution of mercuric chloride is a poor conductor of

MERCURY 197

electricity; it is dissociated to a slight extent only and acts quien
differently in many cases from a solution of the nitrate, which is a good
conductor of electricity and therefore contains a good many mercuric
ions. The cyanide differs from the nitrate even more, as we shall see.

Mercurie bromide is difficultly soluble in water (94 cc. of water
at 9° dissolve only 1 gm. of the bromide), but is readily soluble in alcohol,
and still more soluble in ether. The iodide is more difficultly soluble.

The halogen compounds of mercury readily form complex com-
pounds with the halogen compounds of the alkalies, which are very
stable.

Mercury compounds are furthermore characterized by the readi-
ness with which they undergo hydrolysis, forming insoluble basic
salts. Thus the sulfate is decomposed when diluted largely with
water (particularly on warming) into a yellow insoluble basic salt:

3HeSO1+2H20 — 2H2S014-+Hg30280x.

The presence of hydrogen ions prevents this hydrolysis.
The nitrate also is readily hydrolyzed into more or less insoluble
basic salts, according to the dilution.

Hg(NO3)2+H20 @ HNO3+He(OH)NOs,
or, |
2Hg(NO3)2+2H20 — Hg20 (OH)NO3+3HNOs.

REACTIONS IN THE WET WAY

A solution of mercuric chloride and one of mercuric nitrate are
used for these reactions.
1. Potassium Hydroxide precipitates yellow mercuric oxide:

HgCle +20H — 2Cl->-+H20 +HgO.

The hydroxides of the noble metals are exceedingly unstable;
they lose water, as a rule, even in aqueous solution, forming the anhy-
drous oxide.

On adding a lesser amount of caustic potash to a solution of mercuric
chloride, a reddish-brown precipitate of basic chloride is obtained:

2HeClo+20H- > 2CI-+H20+Hg20Ch,
or
3HeClo-+40H — 4CI-+2H20+Hg302Ch.

198 HYDROGEN SULFIDE GROUP

Pe Mercuric oxide and the basic salts are readily soluble in acids.
2. Ammonia produces ir a solution of mercuric chloride a white

precipitate of mercurjgaminochloride:’
HgCle+2NH3 — NH,zt+ CE +Hg (N H2)Cl.

This compound, the so-called “ infusible precipitate,” volatilizes
before it melts. Itris soluble in acids, and in hot ammonium chloride,
forming the “ fusible precipitate ”’

Hg(NH2)Cl+NH4Cl=Hg(NH3)2Cle.

If ammonia is allowed to act upon mercuric nitrate a white oxy-
amino compound is always formed: .

| Tues
2Hg(NO3)2+4NH3+ H20 — 83NH4NO3+0 NH2-NOs.

3. Potassium Iodide produces a red precipitate of mercuric iodide,
HgCle+2I- — 2Cl"+Hglz,
soluble in excess iodide ions, forming a colorless complex anion:

Hgl2+21° > [Hela].

This complex anion is scarcely dissociated at all into simple mercurie
cations, for the solution gives no precipitate with caustic soda or potash.
The alkaline solution is the so-called “‘ Nessler’s reagent,’’ and serves for
the detection of very slight traces of ammonia. Erte is formed in this

reaction the brown-colored compound, o< “8 Nv I, which is
Hee

soluble in an excess of the “ Nessler’s reagent,’ with an intense yellow
color (cf. p. 90).

4. Alkali Carbonates precipitate from both the chloride and the
nitrate a reddish-brown basic carbonate in the cold,

AHgClo+4Na2CO03 =8NaCl+3C02 T +Hg.03-COsz,

which on boiling loses carbon dioxide and is changed into yellow mer-
curic oxide.

5. Alkali Bicarbonates produce no precipitation in a solution of
mercuric chloride, but do cause precipitation from mercuric nitrate:

4H g(NOz3)2+8NaHCO;=8NaN0O3+4H20+7CO2 T +Hgs03COs.

6. Hydrogen Sulfide produces in solutions of mercuric salts a
precipitate which is at first white, then yellow, brown, and finally

. >

MERCURY 499

black. The white precipitate is formed according to the followin
equation:
3HgClo+2H2S > 4HCl+Hg;ClSs.

By the further action of hydrogen sulfide, black mercuric sulfide is
finally obtained:
HgsCloS2+ Hes =2HCl 4 3H es.

Mercuric sulfide is insoluble in dilute boiling acids. Hot concen-
trated nitric acid transforms it gradually into white Hg3S2(NOs)2,

9HgS+8HNO3=2NO 7 +3S+4H20 + 3Hg3S2(NO3)o,

which by long boiling is changed into the soluble nitrate.
It dissolves readily in aqua regia, forming the chloride with separa-
tion of sulfur:

3HgS+6HCl1+2HNO3=3HgCle+388+2NO 7 +4H20.

Mercuric sulfide is insoluble in caustic soda and potash solutions,
and in ammonium sulfide, but it dissolves readily in sodium or potas-
sium sulfide:

HgS+K2S = Hg(SK)po.

By dilution with water this compound is completely hydrolyzed
into mercuric sulfide, potassium hydrosulfide and potassium hydroxide:

He(SK)2-+H20 @ KOH-+-KSH+HgS.

Therefore it is always necessary to dissolve the mercuric sulfide
with considerable potassium sulfide, or with little potassium sulfide
and considerable caustic potash, in order to prevent this hydrolysis:

The fact that Hg(SK)2 is so readily hydrolyzed explains the forma-
tion of cinnabar in nature. In the interior of the earth the thio
compound is formed, which is brought by springs to the surface and there
undergoes the above decomposition.
7. Potassium Cyanide produces in a solution of ‘mercuric chlo-
ride no precipitation, because the cyanide, as well as the chlo-
ride, forms readily soluble complex compounds with alkali chlorides.
The following are known: K[HgCls], Ko[HgCls], K{Hg(CN)eCl],
Ke[Hg(CN)2Cle] and Ke[Hg(CN)4]. »

In a concentrated solution of merctiric nitrate, potassium cyanide
produces a precipitate of mercuric cyanide, soluble in considerable
water and jn potassium cyanide:

Hg(NO3)2+2CN — 2NO3 +Hg(CN)>2.

200 HYDROGEN SULFIDE GROUP

Mercuric cyanide is the only cyanide of the heavy metals that is
soluble in water. It dissolves mercuric oxide perceptibly, forming
the complex compound (HgCN):O- Mercurie cyanide is not pre-—
cipitated by alkali carbonates or by caustic alkalies, because the
mercuric oxide is soluble in mercuric cyanide. It is not decomposed
by dilute sulfuric acid, although it is by the halogen acids—most
difficultly by hydrochloric acid, and most readily by hydriodic acid;
hydrogen sulfide decomposes it with precipitation of mercuric sulfide:

Hg(CN)2+H2S8 = 2HCN-+H¢s.

8. Neutral Alkali Chromates precipitate yellow mercuric chromate
from both the chloride and nitrate solutions. On long standing or
by boiling, the precipitate becomes red, a basic salt being probably
formed.

9. Alkali Dichromates throw down a yellowish-brown precipitate
from the nitrate solution, but not from the chloride.

10. Ferrous Sulfate reduces mercuric nitrate on boiling to metallic
mercury:

Hg(NO3)2+2Fet*+ — 2Fe* +*+2NO3°+Heg.

Mercuric chloride and cyanide are not reduced by ferrous sulfate.
11. Stannous Chloride reduces mercuric salts, at first to insoluble
mercurous chloride (calomel),

2HgCle +Snt wae Sn* sail *+HgoClo,

and by further action to metal,
Hge2Cle +Snt + = Snt aa ++9Hg¢.

Metallic mercury separates out in the form of a gray powder.
By decanting the solution, and boiling the residue with dilute hydro-
chloric acid, the mercury appears in tiny globules.

12. Copper, Zinc, and Iron precipitate mercury from solutions
of its salts:

HgCle+Fe = Fe* ++2Cl +Hg,*

HgCle+2Cu = CueClo+Hg.

On placing a drop of mercury solution (whether of a mercurous or a
mercuric salt) upon a piece of bright copper-foil, a gray spot is formed,
which, after being dried, becomes bright as silver on rubbing.

* This reaction is employed for detecting metallic iron in the presence of FeO.
If an excess of HgCl, is present HgeCl, is formed (ef. p. 148).

MERCURY 201

Mercurous Salts

The mercurous salts all contain the bivalent mercurous group
Hg2** and are changed more or less readily into mercuric salts, splitting
off one atom of mercury from the molecule. Mercurous salts contain-
ing oxygen, like mercuric salts, are readily hydrolyzed in dilute aqueous
solutions; thus the nitrate is decomposed according to the equation

He2(NO3)2+HOH = HNO3+Heg2(OH)NOs.

Mercurous chloride (calomel) is insoluble in water and hydro-
chloric acid, but soluble in nitric acid and aqua regia.

REACTIONS IN THE WET WAY

1. Caustic Potash precipitates black mercurous oxide:
Hge (NOsz)2 +20H” — 2NO3 +HeO +Hg20.

2. Ammonia produces a black precipitate of mercuric amino salt
with metallic mercury:

Hg. .
4Hg(NOs)2-+4NHs-+H20=3NHsNO2+0C _DNHLNO: + 2H.
H

It can easily be shown that this precipitate contains metallic mer-
cury by rubbing a piece of pure gold over it; silver-lustrous gold amal-
gam will be formed.

Mercurous chloride gives with ammonia a mercuric amine. with
separation of metallic mercury:

HgeCle + 2NH3 = NH4Cl +Hg (NHo2) Cl +Hg.

By boiling the black precipitate with dilute hydrochloric acid
or with concentrated ammonium chloride solution, the mercuric amine
goes into solution, leaving behind drops of mercury. .
3. Alkali Carbonates give, first, a yellow precipitation of the car-
bonate, which quickly becomes gray, owing to the formation of mer-
curic oxide, metallic mercury, end carbon dioxide:

Hge (NOsz)2 +NaeCO3 = 2NaNO3 +Hge2COs,
and
HgeCO3 = HgO+Hg-+-CO2 7 .

4, Ammonium Carbonate yields the same precipitate as ammonia.
5. Hydrogen Sulfide immediately throws down a black precipitate
of mercuric sulfide and mercury (difference from mercuric salts) :

Hgo(NO3)2+H28 =2HNO3+HgS-+ Hg.

202 HYDROGEN SULFIDE GROUP

The black precipitate does not dissolve completely in potassium sul-
fide, the mercury remaining insoluble, but in alkali polysulfides it
dissolves.

6. Hydrochloric Acid and Soluble Chlorides precipitate white
mercurous chloride (calomel),

Hge (NO3)2+2Cl — 2NO37° + Hg2Clo,

insoluble in water and dilute acids, soluble in strong nitric acid and
aqua regia. On boiling for a long time with water, calomel becomes
gray, owing to a partial decomposition into mercuric chloride and
mercury.

On boiling with concentrated sulfuric acid, mercuric sulfate is
formed with evolution of sulfur dioxide and hydrochloric acid:

(a) Hge2Cl> +He2SO4 = 2HCl+ HgeSO4.
(b) HgeSO4+2H2804 = 2H20+S802 fT +2HgSOx.

7. Neutral Potassium Chromate precipitates red mercurous chromate
on boiling (cf. p. 141):
Hge(NOz)2+ KeCrO4 = 2KNO3+HegeCrO4.
8. Potassium Iodide precipitates green cndctneees iodide,
Hge(NOz)2+2KI =2KNO3+Hgols,

partly soluble in an excess of the Iprecipitant, with the formation of
potassium mercuric iodide and separation of mercury:

HgeIo+21 — [Hgl4]"+Hg.

9. Potassium Cyanide precipitates metallic mercury, mercuric
cyanide being formed at the same time:

Hge(NOs)2+2KCN =2KNO3+He(CN)o+Heg.
10. Stannous Chloride precipitates gray metallic mercury:
Hge(NO3)2+Snt * > Sn*+++*++2Hg¢.

REACTIONS OF MERCURY IN THE DRY WAY

Almost all mercury compounds sublime on being heated in the
closed tube. Mercuric chloride melts first, then vaporizes, forming a
crystalline deposit on the cold sides of the tube. Mercurous chloride
sublimes; the subliraate is almost white, but there is a slight grayish
tint owing to the decomposition of a small part of the substance into

MERCURY | 203

mercuric chloride and mercury. Mercurie iodide yields a yellow sub-
limate, which becomes red on being rubbed with a glass rod. Mercury
compounds containing oxygen (all more or less unstable) yield mercury.

The sulfide gives a black sublimate.

All compounds of mercury, when mixed with sodium carbonate
and heated in a closed tube, yield a gray mirror, consisting of small
globules of mercury. In order to make the drops more apparent,
place a piece of filter-paper over a glass rod, and rub the mirror
with it. The small drops thenj run together into large ones, stick
to the paper, and can be removed from the glass.

Detection of Mercury in Urine*

Treat 500 to 1000 cc. of urine in a beaker with 0.5 per cent hydrochloric or
sulfuric acid, add 0.5 gm. of brass wool (such as is used for the ornamentation
of Christmas trees) and, while heating on the water-bath to 60° or 80°, pass air
through the liquid for from ten to fifteen minutes, to keep it in constant motion.
Any mercury present is replaced by copper: HgCl.+Cu =CuCl.+Hg.

The mercury, as fast as it is set free, amalgamates with the excess of copper
present. Pour off the liquid from the tiny threads of brass and wash thoroughly
by decantation-with distilled water, then with alcohol, and finally with ether.
Press the brass thread between layers of filter-paper, to free them from any
adhering ether, and roll them between the fingers into a small pellet. Intro-
duce this pellet into a thoroughly cleansed and perfectly dry glass tube, 10 em.
long, 0.5 em. wide and closed at one end. With the aid of the blast flame,
draw out a capillary of about 1 mm. width in the tube about 0.5 cm. away from
the brass, toward the open end of the tube. After cooling the tube, heat the
bottom of it, in which the sample rests, to dark redness. This causes the mer-
cury to distill off and it is condensed in the colder portion of the tube in a

gray mirror consisting of tiny drops. If any considerable amount of mercury is
present, e.g., more than 1 mg., the drops of mercury ean be distinctly seen with
alens. If less than 1 mg. of mercury is present in the urine, it is very difficult
to distinguish the mirror. In this case to make it perceptible, transform the
mercury into scarlet-red mercuric iodide. To accomplish this, place a small
crystal or two of iodine in a test-tube and cut off the tube containing the mer-
cury mirror just above the place where the ball of brass rests and place the
part of the tube containing the mercury in the test-tube. Cautiously heat the
bottom of the test-tube over a gas flame. As soon as the violet vapors of iodine
reach the place where the mercvry was deposited, the latter is transformed, by
very gentle heating, into the red iodide, which can be seen most distinctly by
removing the little tube and laying it upon a piece of white paper. This method
is very sensitive and permits the positive recognition of as little as 0.4 mg. of
HgCl..f i

*P. Fisrinaer, Z. anal. Chem. (1888), 27, 526.

+ For other methods of detecting mercury in urine, see JOLLES, Z. anal. Chem.,
39, 230 (1900), Merart, J. Pharm. Chim. [5] 19, 444 (1889); and OprpENHEIM,
Z. anal. Chem., 42 (1903), 431.

204. HYDROGEN SULFIDE GROUP

Detection of Mercury Vapors in the Air

Place a piece of pure gold leaf in a small glass tube and draw the air to be
tested through the tube for an hour, at a rate not greater than one liter per
minute. Meanwhile evacuate a Geissler tube, of the form shown in Fig. 19,
by means of a water pump (not a mercury. pump!) and finally close both stop-
cocks. Place the gold leaf, which now contains as
amalgam any mercury that was present in the air
tested, in the tube » and suddenly open the cock a,
which has a wide bore; this causes the gold to be
sucked into the tube, stopping at c, the mouth of the
capillary opening. The next step is to replace the air
in the tube by hydrogen. Introduce hydrogen gas,
obtained from a Kipp generator and dried by concen-
trated sulfuric acid, at a@ and allow the gas to pass out
at b. After a rapid stream of the gas has passed
through the tube for three minutes, close the cocks a
and b, without disconnecting the Kipp generator, con-
nect b with the suction pump* and evacuate the
apparatus for a minute or two; then close b, and open
a (which causes more hydrogen to enter the apparatus);
close a again, open b and once more evacuate the
apparatus. Repeat this alternate introduction otf
hydrogen gas and evacuation five or six times. In
this way the air is entirely replaced by hydrogen. .
Finally evacuate the tube for five or ten minutes and
close the cock b. Place the capillary in front of the slit.
of a spectroscope and allow the secondary current of
an induction apparatus to pass through the tube. In
the presence of the merest trace of mercury, the
characteristic green line 546 uy is distinctly visible in
the cold, and with somewhat larger amounts of mercury
the indigo-blue line at 456 uy can be seen. If the wad
of gold leaf is cautiously warmed with the Bunsea
flame, the mercury spectrum appears still more

Fia. 19. sharply.

Remark. This test is so extremely sensitive that a
blank test performed in places where work with mercury has been performed
will often show the presence of this element in the atmosphere.

If the apparatus has been once used for the detection of mercury it must be
thoroughly cleansed before it is used again for this purpose. To this end,
remove the gold and allow aqua regia to remain in the tube for several minutes.
Draw out the acid and rinse the tube three times with distilled water, once with
absolute alecoholt and finally dry by passing dry hydrogen through the tube
for five minutes, while warming it at the same time. Ignite the gold gently to

* Between the water suction-pump and the Geissler tube, a calcium chloride
drying tube should be introduced.

+ All these operations must be carried out in a space where there are positively
no mercury vapors present in the atmosphere.

LEAD 205

distill off any mercury it contains. If now on introducing the gold and evacuat-
ing the apparatus, the mercury spectrum is no longer visible, the tube is ready
for a new experiment.

It may be mentioned that the two platinum wires in the Geissler tube must
not be provided with aluminium points, because aluminium amalgamates with
mercury, and when the points are once amalgamated it is impossible to free
the tube sufficiently from mercury to permit its use for subsequent experiments.

LEAD, Pb. At. Wt. 207.1
Sp. Gr. =11.36-11.39. M. Pt.=327.4°. B. Pt.=1600°

Occurrence.—Galena, PbS, isometric; cerussite, PbCOs3, ortho-
rhombic and isomorphous with aragonite, CaCOs3; anglesite, PbSOu,
orthorhombic, isomorphous with anhydrite, CaSO., celestite, SrSOq,
and barite, BaSO4; pyromorphite, Pbs(PO4)3Cl, hexagonal; mimetite,
~Pbs(AsO4)3Cl; vanadinite, Pbs(VO4)3Cl. The last three minerals are
isomorphous and belong to the apatite group. Other minerals which
may be mentioned are wulfenite, PbMoO,, tetragonal, isomorphous
with stolzite, PbWO., and the monoclinic crocoite, PbCrO..

Lead is a bluish-gray metal. It is attacked by all acids. As,
however, most lead salts are difficultly soluble in water, it usually
becomes coated with a layer of salt, which protects it from further action
of the acid. Thus lead is immediately attacked by dilute sulfuric
acid according to the equation

Pb-+H2S01=PbSO04+Hp ¢ .

But, as lead sulfate is insoluble in dilute sulfuric acid, the reaction
quickly ceases. Upon this principle rests the use of ‘‘ lead chambers ”’
in the manufacture of sulfuric acid, and the use of “ lead pans” for
the concentration of the dilute “‘ chamber acid.”’ It has, however, been
found from experience that the sulfuric acid should not be concentrated
too much in lead pans—stopping when a 78-82 per cent acid is obtained.
The protecting layer of lead sulfate is soluble in hot concentrated sul-
furic acid, forming soluble lead bisulfate, PbSO4-+ H2SO4 —> PbH2(SO4)e,
so that the hot concentrated acid can act on the freshly-exposed surface
of lead:
Pb+3H2804 — 2H20+PbHe(SO4)2+S802 TF.

Lead behaves quite similarly on treatment with hydrochloric acid.
On the surface a protecting coating of lead chloride is obtained, which
is soluble in hot concentrated hydrochloric acid, forming HPbCls.
Lead is soluble, therefore, in concentrated hydrochloric acid.

Pb+3HCl=HPbCl3+He f .

206 HYDROGEN SULFIDE GROUP

Hydrofluorie acid attacks lead similarly, forming a_ protecting
layer of lead fluoride, which is insoluble in hydrofluoric acid. Con-
sequently lead retorts can be used for the distillation of hydrofluoric |
acid and in the preparation of hydrofluoric acid by means of fluorite
and sulfuric acid.

Nitric acid is the proper solvent for lead. Lead nitrate is insoluble
in strong nitric acid, so that lead does not dissolve in concentrated nitrie
acid; the solution must be sufficiently dilute to prevent the separation
of the lead nitrate formed.

Lead forms the following oxides: lead suboxide, Pb2O; lead oxide
(litharge), PbO; lead sesquioxide, Pb2O3; minium (red lead), Pb3O4;
lead peroxide, PbOs. oe

Of these oxides, PbO alone is the anhydride of a base;* from it
the salts of lead are derived, in which the lead is bivalent. This lead
monoxide (litharge) is a yellow powder, which melts at about 980° C.,
and solidifies on slow cooling, forming tetragonal crystals (needles).
It is slightly soluble in water with an alkaline reaction, and is readily
soluble in dilute nitric acid. :

Lead suboxide, Pb2O, is formed as a black velvety powder on heat-
ing the oxalate to about 300° C.:

2PbC204=3CO2 T +CO FT +Pb20.

On heating the suboxide in the air, it becomes readily oxidized to
lead monoxide.

Lead dioxide, PbO2z, must be considered as the anhycride of ortho-
plumbic acid, H4PbO,, or metaplumbic acid, H2PbOs,

OH
“OH tug
Pb OH oF Pb=0 ,
OH NOH

Orthoplumbic acid Metaplumbic acid

just as $102, SnOz, MnOz, are anhydrides of silicic, stannic, carbonic,
and manganous acids. The acid H2PbOs3 is formed by the oxidation
of lead hydroxide, Pb(OH)s, in alkaline solution by means of hypo-_
chlorites, chlorine, bromine, hydrogen peroxide, or potassium persul-

fate:
Pb(OH)2+20H-+Cle — H20+2Cl-+He2PbOs3.

The brown metaplumbic acid which separates out goes over at
100° C. into the anhydride; and the latter on ignition loses oxygen,
changing into yellow lead monoxide. The other two oxides of lead,

* Although Pb(C,H;02), is known.

. LEAD 207

Pb203 and Pb3O4, may be regarded as salts of the plumbic acids;

Pb203 as lead metaplumbate, PbPbO3, and Pb304 as the lead ortho-
plumbate, Pb2PbOx.

, Pb2O03 is obtained as a yellow precipitate on gently oxidizing an

alkaline solution of lead monoxide by means of hypochlorites, halo-

gens, hydrogen peroxide, or persulfates,

2Pb(OH)2+20H +Cle — 2Cl-+3H20+Pb203,
and the red minium, Pb30.4, by igniting lead oxide or lead carbonate
for some time in the air at about 430° C.:
3PbO0+0=Pb30..

Both Pb2O03 and Pb304 behave chemically as salts; for, on treating
with nitric acid, brown plumbic acid and lead nitrate are formed, which
corresponds to the action of nitric acid on, say, lead carbonate:

PbPbO3 +2H+t— Pb+++HoPbOs,
PbsPbO.+4H*t = 2Pb+ ++ HO +HoPbOs.

These salt-like oxides * are perfectly analogous to those of mangan-
ese; on treatment with hydrochloric acid they yield chlorine, the
plumbic acid, at first set free, behaving like a peroxide:

PbO2+4HCl=2H20+PbCle +Cle J ;
Pbe03+6HCl=3H20+2PbCle+ Clo T ;
Pb304+8HCl=4H20+3PbCle+Cle fF .

REACTIONS IN THE WET WAY

Most lead salts are difficultly soluble or insoluble in water; but all
dissolve in dilute nitric acid, excepting, perhaps, fused lead chro-
mate, which is very difficultly soluble.

1. Potassium and Sodium Hydroxides precipitate white lead hy-

droxide,
| Pb**+20H- — Pb(OH)s,

soluble in an excess of the precipitant, forming a plumbite:
Pb(OH)2+OH — H20+HPbOz.

Pb(OH)s is also slightly soluble in water which is free from carbonic
acid. The aqueous solution of lead hydroxide is slightly alkaline.

* Besides the lead salts of plumbic acid, alkali and alkaline earth salts are known.

208 HYDROGEN SULFIDE GROUP

2. Ammonia precipitates the white hydroxide, insoluble in excess
of the reagent.

3. Alkali Carbonates precipitate white basic lead carbonate. Alkali
bicarbonates precipitate the normal carbonate.

4. Sodium Phosphate precipitates white lead phosphate,

38Pb**+4HPO4" — 2H2PO4-+Pb3(PO4)o,

insoluble in acetic acid, readily soluble in nitric acid, caustie soda
or potash.

5. Potassium Cyanide precipitates white lead cyanide, insoluble
in an excess.

6. Hydrochloric Acid or Soluble Chlorides precipitate from moder-
ately concentrated solutions flocculent, white lead chloride:

Pb*++2Cl- > PbCh,

difficultly soluble in cold water (135 parts of water dissolve 1 part
of PbCl2), but much more soluble in hot water; .on cooling the solution,
lead chloride separates in the form of glistening needles or plates.
Lead chloride is much more soluble in concentrated hydrochloric acid
and in a concentrated solution of a chloride of an alkali than it is in
water, as it forms complex compounds with these substances such as
HPbCls, KPbCl3, which are, however, decomposed on dilution with
water, with separation of lead chloride.
7. Potassium Iodide precipitates yellow lead iodide:

Pb*t+2I- —> Pblo.

The iodide is much less soluble in water than the chloride (195 ce.
of boiling water dissolve only 1 gm. of lead iodide), forming a colorless
solution from which lead iodide separates on cooling, in*the form of
gold-yellow plates.

The iodide dissolves to a considerable extent in hydriodie acid,
and in a solution of an alkali iodide, forming lead hydriodie acid;
HPblIs or one of its salts (such as KPbI3), all of which are decomposed
on dilution, with deposition of lead iodide. 7

8. Alkali Chromates produce a yellow precipitate of lead chromate,

Pbtt+Cr0O)>:* PbCrO,4
and

2Pb* +4 Cr.07-+2C2H302 +H20 — 2HC2H302+2PbCrO,.

Lead chromate is insoluble in acetic acid, but soluble in nitrie acid
and in caustic alkali.

9. Hydrogen Sulfide produces in dilute lead solutions (from

LEAD © 209

slightly acid solutions, as well as from neutral or alkaline ones) a black
precipitate of lead sulfide:

Pb*t++H2S = 2Ht-+PbS.

From hydrochloric acid solutions an orange-red precipitate of lead
sulfochloride is at first obtained,

2PbCle + HS aR 2HCl a Pb 2CloS,

which is decomposed immediately by more hydrogen sulfide, forming
the black lead sulfide. In this respect lead salts are similar to mercuric
salts (see p. 198).

Lead sulfide is soluble in dilute, boiling, 2-normal nitric acid, form-
ing lead nitrate, with separation of sulfur: :

3PbS+2HNO3+6H*— 3Pbtt+4H2,0+2NO0-+38S.

The reaction usually goes a little further; some of the sulfur is
oxidized to sulfuric acid, forming insoluble lead sulfate. ‘The amount
of sulfuric acid formed (and therefore of the lead sulfate also) increases
with the concentration of the acid.

Lead sulfide is also soluble in strong hydrochloric acid:

PbS+2H*t = Pbht*++H.S 7.

10. Sulfuric Acid and Soluble Sulfates cause in solutions of lead
salts the separation of white, difficultly soluble lead sulfate:

Pb* i HeSO4 — 2H +PbSOxz.

One part of the salt dissolves at the ordinary temperature in 22,800 ce.
of water; in water containing a little sulfuric acid it is still less soluble,
while in alcohol it is insoluble. Lead sulfate dissolves perceptibly
in hot, concentrated acids, forming Pb(HSO4)2. On cooling the hydro-
chlorie acid solution, lead chloride separates out in needles. Almost
all the sulfuric acid of commerce contains some dissolved lead sul-
fate. In order to detect this, 200-300 cc. of the concentrated acid
should be diluted with an equal volume of water and allowed to stand
twelve hours, whereby the lead sulfate separates as a white powder.

Besides bhing soluble in acids, lead sulfate is easily soluble in
eaustic alkalies, and in solutions containing the ammonium salts
of many organic acids. This last property is of great importance for
the analytical chemist, as it offers a means for separating lead sulfate
from barium sulfate, silica, etc., which remain undissolved. Ammo-
nium acetate or ammonium tartrate is usually used as the solvent.

The reason lead sulfate dissolves in a concentrated solution of

210 HYDROGEN SULFIDE GROUP

ammonium acetate is due to the formation of lead acetate, which is
ionized to only a very slight extent in the presence of an excess of acetate
ions:

PbSO4+2C2H302, — Pb(C2H302)2+804> +2H20.

From this solution the lead can be precipitated as chromate by the
addition of potassium chromate, as sulfate upon the addition of dilute
sulfuric acid or as sulfide by ammonium sulfide. 3

Similarly, ammonium tartrate dissolves lead sulfate by forming
a tartrate which does not ionize to any extent into simple lead cations.

REACTIONS IN THE DRY WAY

Heated with sodium carbonate on charcoal, all lead compounds
yield a malleable button, surrounded with an incrustation of the
yellow oxide. On the oesctal stick also, the malleable button is
readily obtained.

Lead glass turns black on heating in the reducing flame, owing
to the separation of lead.

BISMUTH, Bi. At. Wt. 208.0
Sp. Gr.=9.8. M. Pt.=270°. B. Pt. about 1435°.

Occurrence.—Bismuth usually occurs native with nickel and co-
balt ores. The following ores are of no great importance: Bis-
mite, BigOs; bismuthinite, BisS3; emplectite, BizSsCuz; bismutite,
3[CO3][BiOH]-5Bi(OH)s.

Bismuth is a brittle, reddish-white metal which crystallizes in the
hexagonal system. The proper solvent for bismuth (as is the case
with most other metals) is nitric acid. Hydrochloric acid does not
attack bismuth, and sulfuric acid dissolves it only on warming.

Bismuth forms three oxides: bismuth trioxide, BizO3, bismuth
tetroxide, BizO4, and bismuth peroxide, BizQ5/ |

Bismuth trioxide is a basic anhydride,* from which the salts are
derived. Bismuth pentoxide, a brownish substance, acts as an acid
anhydride, forming an acid, HBiOs, corresponding to metaphosphoric
acid. Salts of this acid have never been prepared in a pure state.
On igniting, BizOs loses oxygen, forming yellow BisQ3. It dissolves
in hydrochloric acid with evolution of chlorine, forming a salt of
trivalent bismuth:

Bi.0;+10HCl=5H20+2BiCls+2C lo.

* Bismuth trioxide acts as a weak acid under some circumstances (ef. foot-note,
p. 211},

BISMUTH 211

Bismuth tetroxide is a brown powder which is sometimes used as an
efficient oxidizing agent. Commercial sodium bismuthate is probably
a mixture of NaBiO3 and BigQOx. .

Bismuth salts are mostly colorless, and are all insoluble in consider-
able water, on account of being hydrolyzed into an insoluble basic
salt; thus the chloride is CuarrtAuyely decomposed into bismuth
oxychloride,

BiCls3 +H20 = 2HCI1+BiOCl,

insoluble in tartaric acid (difference from antimony).

Bismuth oxychloride is readily soluble in hydrochloric acid, the
above equation taking place from right to left. The reaction, there-
fore, is reversible and the relative amounts of water and hydrochloric
acid present determine in which direction the reaction will go. On
adding water to a slightly acid solution of BiCl3, a white precipitate
of the oxychloride appears immediately. On carefully adding hydro-
_ chloric acid, the precipitate again dissolves, but may be reprecipitated
by the addition of more water. All the other compounds of bismuth
act as the chloride. The nitrate yields, at first, an amorphous pre-
cipitate of BiONOs,

Bi(NO3)3+H20 @ 2HNO3+BiO(NOs3),
which becomes more basic on further addition of water, and crystalline:
2BiO(NO3) +H20 = BizO2(OH) (NOs) +HNOs.

This last compound is the bismuth subnitrate which is so much used
in medicine.
REACTIONS IN THE WET WAY

:. Potassium Byarasise precipitates, in the cold, al bismuth

hydroxide,
Bit t+++30H- — Bi(OH)s,

which, on boiling, becomes pale yellow:
Bi(OH)3 — H20+Bi0O(OH).

Both of these hydroxides are insoluble in an excess of the pre-
cipitant,* but are readily soluble in acids.
On adding to the alkaline solution, in which the hydroxide is sus-

*In very concentrated KOH, Bi(OH); dissolves on warming. On cooling,
part of the Bi(OH)s3is precipitated, and on dilution all of it. In this case the
hydroxide acts as a weak acid, like antimony trioxide,

212 HYDROGEN SULFIDE GROUP

pended, chlorine, bromine, hypochlorites, or hydrogen peroxide, the
white or yellowish precipitate becomes brown, owing to the formation
of bismuthic acid: ;

BiO(OH)+20H +Clz2 — H20+2Cl"+HBiO3.

2. Ammonia precipitates a white basic salt (not the hydroxide),
the composition of which varies with the concentration and with
the temperature.

3. Alkali Carbonates precipitate, according to the temperature
and concentration, a number of basic carbonates; one of which is
formed according to the following equation:

2Bit+++3C037+H20 = 2Bi(OH)(CO3)+COr fF.

4. Sodium Phosphate precipitates the white, granular phosphate,
insoluble in dilute nitric acid, difficultly soluble in hydrochloric:

QHPO,4-+ Bit ty HePO4 +BiPOx,.

5. Potassium Cyanide precipitates the white hydroxide (not the
cyanide). The cyanide is at first formed, but is hydrolyzed:

(a) Bit+*++38CN7- — Bi(CN)s.
(b) Bi(CN)3+3HOH =3HCN-+Bi(OH)s.

6. Potassium Dichromate added in excess precipitates yellow bis-
muthyl dichromate,

Cr2077-++2Bit++4.2H20 2 4H*+ (BiO)2Cr20z,

soluble in mineral acids, insoluble in caustic alkalies (difference from
lead).

7. Hydrogen Sulfide precipitates brown bismuth sulfide,
2Bit +++3H2S — BioS3+6Ht,

insoluble in cold dilute mineral acids and alkaline sulfides, soluble
in hot dilute nitric acid, and in boiling, concentrated hydrochloric
acid.

8. Alkali Stannites (an alkaline solution of stannous chloride)
cause a black precipitation of metallic bismuth.* This very sensitive
reaction is performed as follows: To a few drops of stannous chloride,
add caustic alkali until the white precipitate at first produced dissolves
clear. Add this sodium stannite solution to the cold bismuth solution;

* VaNnINO and TREUBERT, Ber., 1898, 1113.

BISMUTH 213

on shaking, a black precipitate immediately appears. The following
reactions take place in this test:

Sn? *+20H™ — Sn(OH)>;
Sn(OH)2+20H — 2H20+Sn02";
3Sn02-+2Bitt*+60H- — 3H20+38n037+ 2Bi.

In making this test, a too concentrated caustic alkali solution should
be avoided and the solution must be kept cold, otherwise the stannite
itself may give a black precipitate.

If too much caustic potash is used, metallic tin will separate out
(ef. p. 170):

2Sn02~+H20 — SnO03°-+20H+Sn.

If too little caustic potash is used, black stannous oxide will be
thrown down in the cold, after long standing; quickly on boiling:

| SnO2"+H20 — 20H-+Sn0.
9. Potassium Iodide precipitates black bismuth iodide,
Bit*+++3I- — Bils,
soluble in excess of the reagent, forming a yellow or orange solution:
Bilgs +I" > [Bila].

By diluting this last solution with not too much water, the black
iodide is reprecipitated, which, on the addition of more water, is changed
into orange-colored basic iodide:

Bil3-++H20 = 2HI-+BiOlI.
10. Metallic Zinc precipitates metallic bismuth:

2Bit +*4+3Zn > 3Zn +++ QBi.

12)
REACTIONS IN THE DRY WAY

Bismuth salts color the non-luminous flame a pale greenish white.
Heated with soda on charcoal before the blowpipe, a brittle button
of the metal is obtained, surrounded by a yellow incrustation of bismuth
oxide.

On heating a compound of bismuth in the upper reducing flame
(p. 68) of the Bunsen burner, the bismuth is reduced to metal,
which is volatilized and~burnt to oxide in the upper oxidizing flame.
On: holding a porcelain evaporating-dish (glazed on the outside and

214 HYDROGEN SULFIDE GROUP

filled with water) just above the oxidizing flame, a barely visible
deposit is obtained, which, on being treated with hydriodie acid,
is changed to scarlet bismuth hydriodic acid:

Bi2O3+8HI =3Hs0+2H[Bily).

The hydriodic acid is easiest obtained by moistening a piece of asbestos,
held in the loop of a platinum wire, in a solution of alcoholic iodine solution
and then setting fire to the moist asbestos. By holding the burning asbestos
under the dish, enough hydriodic acid is formed to change the bismuth oxide
into the red compound. |

By breathing on this deposit, the color disappears, but reappears
as soon as the moisture has evaporated. On exposure to fumes of
ammonia (by blowing the vapors away from the stopper of an ammonia
bottle) the deposit is colored a beautiful orange, owing to the formation
of the ammonium salt of the bismuth hydriodic acid,

H[Bil4] +NHs3 ay NHag|Bil4],

which also becomes invisible on being breathed upon.
By moistening this coating with an alkaline solution of stannous
chloride, black metallic bismuth is deposited.

COPPER, Cu. At. Wt. 63.57
Sp. Gr. =9.84, M. Pt.=1080° C.

Occurrence. Copper occurs as native copper, Cu; cuprite, CusO;
chalcocite, Cu2S; chalcopyrite, CuFeS2; malachite, Cue(OH)2COz3;
azurite, Cus(OH)2(COs)2 and atacamite, Cu2z0(OH)Cl- B20.

Copper is a light red, ductile metal.

The proper solvent for copper is nitric acid:

8Cu+8HNO3 — 3Cut *+6NO3 +4H20+2NO f.

Bright copper is not dissolved by hydrochloric acid alone, but
in the presence of a weak oxidizing agent, e.g., ferric chloride, the
solution of the metal is easily effected. Hot hydrobromic acid dissolves —
it with evolution of hydrogen, forming cuprous hydrobromic acid:

2Cu+6HBr — Ha[CueBre]+He 7.

At the beginning of this reaction the solution usually turns dark
violet on account of the formation of the cupric salt of cuprous hydro-.
bromic acid, owing to the copper being somewhat oxidized on the sur-
face. In this case, however, the solution soon becomes colorless,

COPPER 215°

owing to the reduction of the cupric salt by metallic copper. On adding
water to the clear solution cuprous bromide is precipitated:

{CusBre]~ ~ —> CueBre+4Br-.

Copper is not attacked by dilute sulfuric acid, but it dissolves in
hot concentrated sulfuric acid, forming cupric sulfate with evolution
of sulfur dioxide:

Cu+2H2504 — CuSO4+H20+S802 fT.

The behavior of copper toward acids can be understood by reference to
the electromotive series (p. 41). As copper is below hydrogen in the series it
can be oxidized by hydrogen ions only when the concentration of cupric ions
is kept very low (cf. p. 43). Hydrobromic acid dissolves copper because the
slightly ionized complex is formed. Sulfuric acid dissolves copper by virtue
of the oxidizing power of the hexavalent sulfur.

Copper forms two oxides: red cuprous oxide, CueO, and black
cupric oxide, CuO.

Both oxides are basic anhydrides, forming cuprous and cupric
salts. Salts of the cuprous series contain the bivalent cuprous group,
Cugt+, while proee of tke cupric series contain the simple, bivalent
copper atom Cut *

Copper is also known in the trivalent condition.* If a nitric acid solu-
tion of tellurous acid is evaporated to dryness with a little copper nitrate
and the residue is treated with KOH solution (1 : 5) it dissolves. If to the clear
solution 4 to 6 gms. of (NH4)2(SOx)2 are added, little by little, while the solu-
tion is at the temperature of the water bath, it becomes pink and the tellurium
is present for the most part as telluric acid but to some extent as the potas-

oes
sium salt of telluro-cupric acid : K[O-Cu TeO,].
N\oZ

A. Cuprous Compounds

The cuprous compounds are extremely unstable, being oxidized
quickly to cupric compounds. The only known cuprous salts are
those with the halogens, the very unstable sulfate and the sulfite.
Cuprous salts are colorless, insoluble in water, but readily soluble in
concentrated halogen acids, forming colorless solutions. Such solu-
tions contain the unstable cuprous halogen acids, probably of the
formula H4[Cuz2X¢6), in which ‘‘ X’’ is either chlorine, bromine, or
iodine. Salts are known which are derived from these acids, e.g.,
K4[CueCle].

* Cf. Moser, Z. anorg. Chem., 54 (1907), 119 and Braunrr and Kuzma, Ber.,
(1907), 3362.

216 HYDROGEN SULFIDE GROUP

The cuprous’ halogen acids are changed dark on contact with air.
The chloride becomes brownish black; the bromide, dark violet; —
probably due to the formation of cupric salts of the cuprous halogen
acids.

The behavior of the cuprous halogen acids toward carbon mon-
oxide is very important; the latter is readily absorbed, forming an
unstable compound:

CueCle+2CO+2H20 @ CuzCle:-2CO:-2H20.

By boiling the solution the compound is decomposed into cuprous
chloride and carbon monoxide; cuprous chloride is used in gas analysis
for the absorption of this gas.

REACTIONS IN THE WET WAY

A solution of cuprous chloride in hydrochloric acid should be used,
which may be prepared as follows: Dissolve 2 gms. of cupric oxide
in 25 ec. of 6-normal hydrochloric acid, pour the solution into a flask —
and add 0.58 gm. of copper filings. Place several copper spirals in the
flask, one end reaching up to its neck, stopper the flask, invert it and —
let it stand several days. The originally dark solution will gradually
become colorless, when it is ready to be used for the following reactions:

1. Potassium Hydroxide produces in the cold a yellow precipitate
of cuprous hydroxide, ,

Cug* ++20H- — Cu2(OH)s,

which loses water at the boiling temperature, changing to red cuprous
oxide:
Cu2(OH)2 =H20+Cu20.
2. Hydrogen Sulfide precipitates black cuprous sulfide,
Cuett+H2S —? 2H*t+Cu.S,

soluble in warm. dilute nitric acid, forming blue cupric nitrate, with
separation of sulfur:

3Cu2S +16HNO3 — 8H20+3S+6Cutt+12NO3-+4NO f.
3. Potassium Cyanide precipitates white cuprous cyanide,
Cuz**+2(CN)~ — Cue(CN)o,
soluble in excess, forming colorless complex cuprocyanide anions: —

Cu2(CN)2+6(CN)” — [Cu2(CN)g] **.

COPPER ? 217

This solution contains no appreciable quantity of cuprous ions, and
gives no precipitation with potassium hydroxide or hydrogen sulfide.
It is estimated that in a normal potassium cyanide solution the ratio of
the concentration of the complex anion to that of simple cuprous ions
is about 1076: 1. This fact is utilized in the separation of copper
from cadmium. ;

In the absence of an excess of CN- ions, however, an appreciable
ionization takes place: [Cue(CNsg)|==—> Cug**++8CN-, and this ioniza-
tion increases as the solution is diluted. From the diluted solutions
the compounds K2Cue(CN)4, K[Cue(CN)s] and finally Cue(CN)e2 are
obtained, which are less complex in nature.

All of these compounds, even in the solid state, are decomposed by
hydrogen sulfide with precipitation of black cuprous sulfide. Con-
sequently, in order to prevent the precipitation of copper by hydrogen sul-
fide, considerable potassium cyanide must be added, more than enough to
form the salt Ke[Cuz(CN)s].

B. Cupric Compounds

Cupric salts are either blue or green in aqueous solution; in the
anhydrous state they are white or yellow.

The chloride, nitrate, sulfate and acetate are soluble in water;
most of the remaining salts are insoluble in water, but readily soluble
in acids.

REACTIONS IN THE WET WAY
A solution of copper sulfate should be used.

1. Potassium Hydroxide produces in the cold a blue precipitate
of cupric hydroxide,

Cut t+20H- — Cu(OH)s,
which on boiling becomes changed into brownish-black cupric oxide.

Cu(OH), is slightly amphoteric in nature and dissolves in very concentrated
KOH or NaOH, particularly on warming, with a blue color. (Cf. p. 180.)

In the presence of tartaric acid, citric acid, and many other organic hydroxy-
compounds, cupric hydroxide is not precipitated by the addition of caustic
alkali, but the solution is colored an intense blue. If this alkaline solution is
treated with d-glucose, aldehydes, arsenious acid or various other substances
having a reducing power, yellow cuprous hydroxide is precipitated from the
warm solution which is changed to red cuprous oxide on boiling. An
alkaline solution of cupric salt containing tartaric acid is commonly used under
the name of Fehling’s solution. It may be prepared by mixing together equal
volumes of a solution containing 34.64 gms. of crystallized copper sulfate in

218 HYDROGEN SULFIDE GROUP

500c¢ c. of water with a solution consisting of 173 gms. Rochelle salt and 52
gms. NaOH in 500 ce. of water. It is best to keep the solutions separate until
they are to be used. Fehling’s solution is a reagent for many kinds of sugar,
aldehydes, hydroxylamine, etc.

2. Ammonia.—On adding ammonia cautiously to the solution
of a cupric salt, a green, powdery precipitate of a basic salt is obtained,
which is extremely soluble in an excess of the wisi forming an azure-
blue solution:

(a) 2CuSO4+2NH.40H = (NH4)2804+Cue(OH) 2804.
(b) Cuz2(OH)2S804+(NH4)2804+6N H3 =2 ((Cu(N Hs) 4|SO4 -H20).

On adding alcohol to the concentrated blue solution, the above
compound is precipitated as a blue-violet crystalline substance, which
gradually loses ammonia on being heated, leaving behind the cupric
salt. On conducting ammonia gas over an anhydrous copper salt,
the ammonia is eagerly absorbed, with the formation of a complex
cupric ammonia salt: CuCl2+6NH3=[Cu(NHs3)6]Cle.

These compounds (which contain as a maximum 6NHg3 to one
atom of copper) are perfectly analogous to the corresponding com-
pounds of nickel, cobalt, and zinc. By the precipitation of the ammo-
niacal solution with alcohol, the compound with 4NHg3 to one atom of
copper is always obtained. ,

[he ionization of the complex cation,

[Cu(NH3)4]*t— Cur ++4NHs,

is slight in the presence of excess ammonia, but much more than that of
the cuprocyanide ion (p. 217).

(3) Hydrogen Sulfide precipitates from neutral or very slightly
acid solutions colloidal, black cupric sulfide; which has a tendency to
form a colloidal solution (p. 58) and run through the filter:

Cu**+H28 — 2H*+CuS.

To prevent the formation of a colloidal solution, the solution must contain
some electrolyte; the hydrochloric acid present when the precipitation is made
is usually sufficient. Another difficulty frequently encountered is due to the
readiness with which a part of the cupric sulfide precipitate is oxidized to sul-
fate by contact with the air. Thus if a filter containing copper sulfide is
allowed to stand in the air, a little cupric sulfate is formed which is soluble in
water. Many cases where the cupric sulfide apparently runs through the
filter are explained in this way. In filtering a copper sulfide precipitate the rule
should be never to let the filter drain completely until the filtration and washing
is over, and the washing should be with dilute Byeey sulfide water, which
serves to prevent any oxidation.

COPPER 219

Copper sulfide is soluble in hot dilute nitric acid, but insoluble
in boiling dilute sulfuric acid (difference from cadmium); it is soluble
in potassium cyanide, forming potassium cuprous cyanide. From a
solution of the latter salt the copper cannot be precipitated by hydrogen

~ gulfide.

- Copper sulfide is appreciably soluble in ammonium sulfide, but is
insoluble in potassium or sodium sulfide* (difference from mercury).
4. Potassium Cyanide produces, at first, yellow cupric cyanide,
which immediately loses dicyanogen, forming white cuprous cyanide.
The latter, as we have already seen, forms soluble potassium cuprous
cyanide with more potassium cyanide:

2Cutt+4CN- — 2Cu(CN)o;
2Cu(CN)2— (CN)2+Cu2(CN)2;
Cuo(CN)2+6(CN)~ — [Cue(CN)s]|>~.

On adding sufficient potassium cyanide to the blue ammoniacal
cupric solution, the complex compound will be decolorized, forming
potassium cuprocyanide, and the reduction of the cupric salt to cuprous
condition in ammoniacal solution is accomplished at the expense of
cyanide ions which are oxidized to cyanate:

2[Cu(NHs)4]*+-+9(CN)- )
| +20H- = [Cus(CNsg)]==+(CNO)~+8NH3+H320.

Hydrogen sulfide will not precipitate cupric sulfide from the color-
less solution of potassium ‘cuprocyanide provided sufficient potassium
cyanide is present (difference from cadmium). Sometimes, when
considerable copper salt is present, the introduction of HeS causes
the formation of a red crystalline precipitate of hydrorubianic acid,
(CSNHe)e. Cf. p. 316.

5. Potassium Thiocyanate, KCNS, precipitates black cupric
sulfocyanate,

Cn+++2(CN8)- — Cu(CNS)>o,

which is gradually changed into white cuprous thiocyanate, or imme-
diately on adding sulfurous acid: |

2Cu(CNS)2+S03°+H20 — 2CNS~ +8047 +2H*+Cu2(CNS)>2.

*In solutions of alkali polysulfides, particularly out of contact with the air,
cupric sulfide dissolves with the formation of compounds of the type NHa4[CuSq]
and K[CuS4]. Cf. Hormann and Hécutien, Ber., 36, 3900 (1903), and Bitz
and Heras, thid., 40, 974 (1907).

220 HYDROGEN SULFIDE GROUP

Cuprous thiocyanate is insoluble in water, dilute hydrochloric
acid, and sulfuric acid.

6. Alkali Xanthates produce in aotaiteans of cupric salts, at first, —
a brownish-black precipitate of cupric xanthate, which splits off dixan- :
thogen, forming finally yellow cuprous xanthate:

S S
sche
NaS-COC2H;+Cut* — 2Nat+Cu(S-COC2Hs)2.

Sodium xanthate

I i ]
2Cu(S:C-OCeHs)2 — Cus(SCOC2H5)2+ (SCOC2H5)2.

Cuprous xanthate Dixanthogen
The reagent, sodium xanthate, is readily obtained by mixing car-
bon disulfide with sodium alcoholate:

|
CS2+Na0C2Hs > NaS-C-OC2Hs.

The alkali xanthates are not used as reagents in testing for copper, but
cupric salts are used in testing for xanthates. The reaction is made use of in
the detection of carbon disulfide in gas mixtures; the gases are allowed to
act upon sodium alcoholate, whereby sodium xanthate is formed if carbon
disulfide is present, and the solution after neutralizing with acetic acid is
tested for xanthate by means of a solution of cupric salt.

7. Potassium Ferrocyanide, K,4{Fe(CN)¢], produces in neutral and >
acid solutions an amorphous precipitate of reddish-brown cupric ferro-
cyanide, ,
2Cutt+[Fe(CN)6]~~ — Cuz[Fe(CN)¢],

insoluble in dilute acids, but soluble in ammonia with a blue color —
(difference from molybdenum ferrocyanide which dissolves in ammonia,
forming a yellow solution). It is also decomposed by potassium
hydroxide: in the cold, light-blue cupric oxide and potassium ferro-
cyanide are formed, while, on warming, black cupric oxide is obtained
(difference from uranium, which yields the yellow uranate both with
ammonia and sodium or potassium hydroxide).

REACTIONS IN THE DRY WAY

The borax, or salt of phosphorus, bead is green in the oxidizing
flame when strongly saturated with the copper salt; blue if containing
only a small amount. The reducing flame decolorizes the bead unless
too much copper is present; in such a case it is reddish brown and
opaque, owing to the separation of copper. ‘Traces of copper may be

CADMIUM 221

determined with certainty as follows: To the slightly bluish bead pro-
duced by the oxidizing flame, add a trace of tin or of a tin compound.
Heat the bead in the oxidizing flame until the tin has completely dis-
solved, then bring it slowly into the reducing flame and finally quickly
remove it. The bead now appears colorless when hot, but ruby-red
and transparent when cold. If, however, the bead is kept too long in
the reducing flame, it remains colorless; but the ruby-red color may be
produced by cautious oxidation. ‘This reaction is very sensitive, and
can also be used for the detection of tin.

Heated with charcoal before the blowpipe (or better still with the
charcoal stick), spongy metal is obtained.

Copper salts color the flame blue or green.

CADMIUM, Cd. At. Wt. 112.4
Sp. Gr.=8.6. M. Pt.=321°. B. Pt.=770°C.

Occurrence.—Cadmium is usually associated with zinc in its ores.
_ It is also found as greenockite, CdS, hexagonal; and as the oxide, CdO,*
isometric. :

The most important commercial salt is the sulfate, 3CdSO4:-8H20.
It is not easily recrystallized. To purify the salt, the concentrated
aqueous solution is treated with alcohol, and the crystals that are
deposited thereby are filtered, washed with alcohol and dried upon
blotting paper.

Cadmium is a silver-white, ductile metal. Heated in the air, it
burns to brown cadmium oxide. The proper solvent for cadmium |
is nitric acid. Dilute hydrochloric and sulfuric acids dissolve it but
slowly, with evolution of hydrogen. Cadmium ‘forms two oxides:
black cadmium suboxide, Cd2O, and brownish-black cadmium oxide,
CdO.

Cadmium suboxide (whose existence is doubted) is formed with
cadmium oxide in small amounts when the metal is burned in the
air. It is also said to be formed, like lead suboxide, by gently
heating the oxalate away from air. There are no cadmium salts
derived from this oxide. Cadmium forms only one series of salts, in
which cadmium is bivalent.

Cadmium salts are mostly colorless, though the sulfide is yellow
or orange. Most of the salts are insoluble in water, but readily soluble
in mineral acids. The chloride, nitrate, and sulfate are soluble in
water.

* With smithsonite in the zine deposits of Monte Poni, Sardinia. Chem. Zig.,
1901, 561.

222 HYDROGEN SULFIDE GROUP

REACTIONS IN THE WET WAY

1. Potassium Hydroxide precipitates white, amorphous cadmium
hydroxide, insoluble in an excess of the reagent (difference from zine
and lead):

| Cd**+20H7- — Cd(OH)>.

On gently igniting the hydroxide the brown oxide is obtained,
which becomes darker on stronger ignition. The ignition of cadmium
nitrate yields the black crystalline oxide.

2. Ammonia also precipitates the white hydroxide, soluble in
excess (difference from lead), forming complex cadmium ammonia
compounds, as with zinc, nickel, etc.

Cd(OH)2+4NHs3 — [Cd(NHs3)4]**F.

In the presence of normal ammonium hydroxide the ratio of the .
concentration of the complex anion to that of the simple cadmium
cation is about 107: 1. In pure water, the ionization takes place to
a much greater extent; by diluting with water and boiling, cadmium
hydroxide is reprecipitated from the solution of the cadmium ammo-
nium compound.

3. Alkali and Ammonium Carbonates precipitate the white basic
carbonate insoluble in excess.

4. Potassium Cyanide precipitates white, amorplfous cadmium
cyanide, readily soluble in excess:

Cd**+2(CN)~ — Cd(CN)po,
Cd(CN)2+2(CN) — [Cd(CN)a]-.

From the solution of cadmium potassium cyanide the above-men-
tioned reagents produce no precipitation. In a normal solution of
potassium cyanide the concentration of the complex anion to that of
the simple cadmium cation is 10!%:1. This is evidently a much
weaker complex than the cuprocyanide anion and for this reason cad-
mium sulfide, though its solubility product is much larger than that of
cupric sulfide, is precipitated by hydrogen sulfide (difference from
copper) :

[Cd(CN)4|>-+H2S — 2CN-+2HCN+CdS.

5. Hydrogen Sulfide produces precipitates varying in color
from a canary-yellow, orange to almost brown, according to the con-
ditions. In neutral solution, whether hot or cold, light-yellow cad-
mium sulfide is obtained in a condition hard to filter. From acid
solutions (containing in 100 cc. from 2 to 10 cc. of conc. H2SO4, or

ARSENIC 223

from 2 to 5 cc. of conc. HCl) yellow precipitates which turn orange in

color are at once thrown down and are easy to filter. The latter pre-
' cipitates are not pure CdS, but always contain more or less Cd2CleS or
Cde(SO4)S. For this reason cadmium should not be determined quan-
titatively as the sulfide.

Cadmium sulfide is insoluble in alkaline sulfides (difference from
arsenic), but is soluble in considerable hydrochloric acid, warm dilute
nitric acid, and hot dilute sulfuric acid (difference from copper).

6. Ammonium Sulfide produces in ammoniacal solutions colloidal
cadmium sulfide, which has a tendency to form colloidal solutions and
pass through the filter. The presence of a concentrated salt solution
prevents its doing this (cf. p. 218)..

REACTIONS IN THE DRY WAY

Cadmium compounds, heated on charcoal with soda, give a brown
incrustation of cadmium oxide.

If a compound of cadmium oxide is reduced in the upper reducing
flame of the Bunsen burner, the cadmium oxide is changed to metal,
which volatilizes, and, in the upper oxidizing flame, goes back to oxide,
which will be deposited as a brown coating if a glazed porcelain dish
filled with water is held just above the flame. This oxide always con-
tains some suboxide mixed with it, and has the property of reducing
silver oxide to metal; so that if the coating of oxide is moistened with
silver nitrate solution a black deposit of metallic silver will be obtained:

Cd.0+2Agt— Cd*+++CdO +2Ag.

This reaction is very sensitive.

If it is desired to test the precipitate produced by Ledoiecn sul-
fide for cadmium in this way, first roast the sample in the oxidizing
flame and then treat it as just described.

ARSENIC, As. At. Wt. 74.96
Sp. Gr. =5.73

Occurrence.—Arsenic is widely distributed in nature, being found
in small amounts in almost all sulfides, as, for example, sphalerite and
pyrites: therefore almost all the zine and sulfuric acid of commerce
contain arsenic.

Arsenic occurs native in kidney-shaped masses; also in the form of
its oxide, As2O3, as isometric arsenolite and orthorhombic claudetite,
it being dimorphous.

294 HYDROGEN SULFIDE GROUP

Mimetite, Pbs(AsO4)3Cl, hexagonal, isomorphous with apatite,
pyromorphite, and vanadinite, is a well-known mineral containing
arsenic oxide.

The most important sources of arsenic are the sulfides, arsenides,
and sulfo salts: realgar, AseSe, monoclinic; orpiment, As2S3, mono-
clinic; arsenopyrite, FeAsS, orthorhombic; niccolite, NiAs, hexagonal;
léllingite, FeAse, orthorhombic; smaltite, (Co,Ni,Fe,)Ase, isometric;
and proustite, As(SAg)3, rhombohedral.

Metallic arsenic is a steel-gray, brittle substance. On being heated
it sublimes, giving off a characteristic garlic odor. The merest trace
of arsenic may be recognized by this odor. The molecule of arsenic
contains, like phosphorus, four atoms, (Asg4). :

Arsenic is insoluble in hydrochloric acid, but readily soluble in
nitric acid and in aqua regia.

Dilute nitric acid dissolves arsenic, forming arsenious acid:

Ass +4HNO3+4H2C = 41H3As03+4NO 7.

Concentrated nitric acid and aqua regia dissolve it, forming arsenic
acid:
3As4 + 20HNO3+8H20 = 12H3As0O4+20NO 7 .

Arsenic belongs to the same natural group of eldnents as nitrogen
and phosphorus, and forms, as they do, two oxides, AsgO3 and AseQs5.

In reality the symbol of the lower oxide is As,Og, but it is customary to use
the simpler symbol, As2,O3.

A. Arsenious Compounds

Arsenic trioxide is formed by the combustion of arsenic in the
air as white, glistening crystals of regular octahedrons. If the vapors
of the trioxide are allowed to cool slowly, they solidify to an amorphous
glass (arsenic glass), which gradually becomes crystalline (white and
opaque, like porcelain).

Arsenic trioxide is known in three different modifications: isometric
arsenic trioxide (white arsenic); monoclinic arsenic trioxide, and
amorphous, glassy arsenic trioxide.

The monoclinic modification is difficultly soluble in water (80 ce.
of cold water dissolve 1 gm. As2O03); while the amorphous, glassy modifi-
cation is much more soluble (25 ec. of cold water dissolve 1 gm. arsenic
trioxide). By treatment of the ordinary modification (white arsenic)
with water, it is not readily wet by the latter; it floats like flour, and
this behavior is very characteristic.

The trioxide dissolves quite readily in hydrochloric acid. particularly

ARSENIC 225
on warming, from which solution it often separates out, on cooling, in a.
beautiful, crystalline, anhydrous condition.
Acting as an acid anhydride it dissolves readily in alkalies, forming
easily soluble arsenites:

As203+60H” — 3H20+ 2AsO3 ;
As203+3C0O3— — 3CO2+2As03°-.

The tri-metal arsenites derived from the ortho acid H;AsO; are usually
unstable. Silver arsenite, Ag;AsO;, is the only well-known salt of this type.
The alkali arsenites are derived from metarsenious acid, HAsO., from pyroar-
senious acid, H,As.O; or from a polyarsenious acid such as HsAsyOs. The only
sodium and potassium salts known are of the types KAsO, and K.H,As,QOq;
of ammonium, (NH,),As,0;. In alkaline solution, however, we may assume
that AsO;~ ions are present.

Free arsenious acid, H3AsO3, has never been isolated; as a very
weak acid it breaks down, like carbonic acid, into water and the anhy-
dride.

Arsenic combines with chlorine directly, like phosphorus, form-
ing the chloride, AsCl3, which behaves exactly like the chloride of
arsenious acid, similar to PCls. It is a colorless liquid, boiling at
134° C., and is decomposed quantitatively, like all acid chlorides,
with water:

AsCl3+3H20 @ 3HCI+H3AsOs3.

The aqueous solution of arsenic trichloride, and the solution of the
trioxide in dilute hydrochloric acid, contain the arsenic as arsenious
acid. As the concentration of the hydrochloric acid increases, the
amount of arsenic trichloride increases, until in very concentrated
hydrochloric acid the arsenic is present almost entirely as trichlo-
ride. By boiling a solution of arsenic trichloride in hydrochloric
acid, arsenic trichloride is given off as a gas. If hydrochloric acid
_is conducted into the solution at the same time (so that the concen-
tration of the hydrochloric acid is kept as large as possible), all the
arsenic can be volatilized from the solution as arsenious chloride.
On evaporating a hydrochloric acid solution of arsenious acid, arsenious
chloride constantly escapes, so that all the arsenic may be volatilized.
If, however, the arsenic is present in the form of arsenic acid, no arsenic
is lost during the evaporation of the solution.

REACTIONS OF ARSENIOUS ACID IN THE WET WAY

The arsenites of the alkalies are soluble in water; the remaining
arsenites are insoluble in water, but soluble in acids.

226 HYDROGEN SULFIDE GROUP

1. Hydrogen Sulfide precipitates from acid solutions yellow, floccu-
lent arsenic trisulfide:

2As(OH)3+3He2S = 6H20+As28s3,
2AsCl3-+3H2S =6HCI+As2Ss.

Arsenious sulfide is insoluble in acids; even boiling 6-normal hydro-
chloric acid does not dissolve it, but by long boiling with 12-normal
hydrochloric acid it is slowly changed to volatile AsCls and HS.
Concentrated nitric acid oxidizes it to arsenic acid and sulfuric acid:

3As2S3+28HNO3+4H20 =9H2801+28NO J +6F3AsO0u.
The sulfide dissolves more readily in ammoniacal hydrogen peroxide:
As2S3+14H202+120H- — 20H20+38047+2AsO4>.

It is also dissolved by alkalies, ammonium carbonate, and alkali
sulfides: |
AseS3+60H- — 3H20+As03 +AsS3 ;

As2S3+3C037 > 83CO2+As03 ¢AsS3_;
AsoSs3 +387 <7 2AsS3.

Just as the anhydride, AseO3, can be referred to the acid, HgAsOs,
so the thioanhydride, AsgSs, can be referred to the thioarsenious
acid, Hs3AsS3, which is not capable of existence in the free state, but is
known in the form of its salts. If one of the latter salts is acidified,
then thioarsenious acid is set free; but it immediately loses Ho,
forming the insoluble thioanhydride:

2AsS3 + 6H* —> 3HeS +AseoS3 is

On treating a mixture of thioarsenite and arsenite with acid, arsenic
trisulfide is also precipitated:

AsO3~ +AsS3 + 6Ht ——y 3H2O +AsoS3 ;

In this last case precipitation is quantitative only when the solution is
dilute; from a concentrated solution H.S escapes, so that more HsS must
be conducted into the solution in order to precipitate all the arsenic.

This property of forming thio-salts accounts for the fact that
hydrogen sulfide produces no precipitation from normal arsenites, —
and only a partial precipitation of AseS3, from mono- and dimetallic
salts: |

‘

ARSENIC 227

AsOz3 + 3He2S —> 3H20+AsS3~ rs
6HAsO37+15H2S — 18H»0+AsoS3-++-4AsS3™
3H2AsO3-+6H2S — 8H20+AsoS3-+-AsSa™.

Consequently, in order to precipitate arsenic completely as tri-
sulfide, it is always necessary that the solution should contain enough
free acid to prevent the formation of soluble sulfo-salts.

2. Silver Nitrate produces in neutral solutions of arsenites a yellow
precipitate of silver orthoarsenite (difference from arsenic acid),

AsO3 +3Ag* — Ag3AsOsz,
soluble in nitric acid and ammonia:
Ag3AsO3+3H* — 3Ag*+H3As0O3.
AgsAsO3-++6NH3 — 3Ag(NH3)2+-+AsOs™.

The first reaction is caused by the formation of nonionized arsenious acid,
which, in the presence of an excess of H™ ions furnishes even less AsO;= ions
than are formed by contact of the very slightly soluble Ag;AsO; in contact
- with water. The solubility in ammonia is due to the fact that the [Ag(NH;).]*
ion in the presence of an excess of NH; furnishes fewer simple Agt cations
than Ag;AsO; in contact with water.

In aqueous solutions of the mono- and dimetallic salts the precipita-
tion is incomplete:

OH
3A OH +3AgNO3=3KNO3+2H3AsO03+Ag3AsO3.

HeAsO3 +3Ag* = Ag3AsO3+2H".

In order to make the precipitation quantitative, an alkali (preferably
ammonia) must be added. As, however, the solution already reacts alkaline,
it is difficult to reach the exact neutral point. Usually too much ammonia is
added. As a rule in qualitative analysis it is unnecessary to accomplish complete
precipitation in this test, as the color of the silver precipitate suffices to show
whether an arsenite or an arsenate is present. To make the precipitation
practically complete, add ammonia drop by drop to a solution of silver nitrate
‘until the precipitate of silver oxide that first forms redissolves; the solution
then contains complex [Ag(NH;).|* cations instead of simple Agt ions. Add
this reagent to the arsenite solution which has been made weakly acid with
nitric acid: |

H;AsO;+3[Ag(NH;).]*+3H* == 6NH,*+Ag-AsO;.

The addition of the nitric acid is necessary, as otherwise the solution will
become ammoniacal, dissolving a part of the silver arsenite.
In case the solution to be tested contains also a chloride, it should be

228 HYDROGEN SULFIDE GROUP

acidified with nitric acid and the chloride precipitated as silver chloride by an
excess of silver nitrate, and filtered off. To the filtrate, dilute ammonia
should be added cautiously. At the neutral zone formed by the ammonia
above the acid solution, a yellow precipitate of silver arsenite will appear.
This reaction is very sensitive.

3. Magnesium Ammonium Chloride produces no _ precipitation
in dilute arsenite solutions in the presence of ammonia (difference
from arsenic acid).

4. Iodine Solution is decolorized by arsenious acid, the latter
being oxidized to arsenic acid:

H,AsO;~+ I, 2 QHt++ 21-+H:AsO,-.

To make the reaction take place quantitatively in the direction left to
right it is necessary to keep the solution neutral; to make the reaction take
place quantitatively in the direction right to left it is necessary to add a con-
siderable excess of hydrogen ions. This is in strict accord with the mass-
action principle. This behavior has been explained by assuming that. free
hydriodic acid is a better reducing agent than iodide ions, but it is more prob-
able that the effect of the acid upon the stability of the arsenic compounds is
more important. In alkaline solutions, the arsenic % more stable in the higher
state of oxidation and, for this reason, the’ arsenite solutions have strong
reducing powers in neutral or alkaline solutions. Arsenious acid is ampho-
teric and forms, as we have seen, a trichloride and trisulfide. In the presence
of an excess of hydrogen ions from some other source, arsenious acid will
not ionize appreciably as an acid and the tendency will be to form Ast*t
cations. Probably arsenic acid is also amphoteric, though to a much less
extent. In strongly acid solutions, the ionization of the arsenic acid is re-
pressed to some extent and there is a tendency to form As***** cations, but
these are far less stable than Ast** cations and, therefore, in strongly acid
solutions an arsenate acts as a vigorous oxidizing agent.

The explanation is in the line with the results obtained in the study of
oxidation potentials (cf. p. 43). The addition of acid decidedly increases
the oxidizing power of arsenic acid, but slightly diminishes the reduction power
of an iodide.

To keep the solution neutral when it is desired to oxidize an arsenite by
means of iodide, it is not advisable to use caustic alkali solution, as this itself
reacts with iodine, forming iodide and hypo-iodite. Sodium bicarbonate is
usually used:

HCO,-+Ht =e H,O+CoO, T .

Equally satisfactory is disodium phosphate, which forms with hydrogen ions
the very slightly ionized H,PO; ions. A normal alkali carbonate can be used
if the solution is saturated with carbonic acid; there are then not enough OH~
ions formed by the hydrolysis of the normal carbonate to react with the iodine.

5. Stannous Chloride (Bettendorff’s Test)—On adding to con-
centrated hydrochloric acid a few drops of an arsenite solution and
then 4 cc. of a saturated solution of stannous chloride in hydrochlorie

ARSENIC 229

acid, the solution quickly becomes brown and then black, owing to the
deposition of metallic arsenic. The reaction takes place more readily
on warming, but a dilute aqueous solution will not give the reaction.
In concentrated hydrochloric acid, however, the arsenic is all present as
trichloride, and this is reduced by the stannous chloride, while arsenious
acid is not:

2As*t*++38n*t — 38ntt++++2As,

B. Compounds of Arsenic Pentoxide

Arsenic pentoxide, which may be obtained by heating arsenic
acid, is a white, fusible substance, and is changed by strong ignition
into arsenic trioxide: AseO5;=As2Os+O2. Arsenic pentoxide is quite
soluble in water, forming arsenic acid:

AseOs +3H2O = 2H3AsQOx4.

Arsenic acid itself may be obtained in the solid state in the form
of orthorhombic prisms corresponding to the formula 2H3AsO4- HO.
At 100° C. water escapes, orthoarsenic acid, H3AsO4, being left behind
as a crystalline powder.

By gentle ignition more water is given off, forming pyroarsenic acid,
H4As207, which on further ignition is changed to metarsenic acid,
HAsOz. In this respect arsenic acid acts exactly like phosphoric acid.
Both the pyro- and the meta-acids readily take on water, and are
changed back to the ortho acid.

The salts of arsenic acid are called arsenates.

As with orthophosphoric acid, mono-, di-, and trimetallic salts are
known: NaHeAsO4, NagHAsO,4 and NagAsOu.

The arsenates of the alkalies are soluble in water; the others are
insoluble in water but easily soluble in acids.

REACTIONS IN THE WET WAY

1. Hydrogen Sulfide on being passed into a cold solution of an
arsenate in 0.3-normal acid does not cause any precipitation until after
a long time, when arsenic trisulfide is formed. If the cold solution con-
tains a large excess of concentrated hydrochloric acid, the arsenic is
precipitated as pentasulfide. If hydrogen sulfide is passed into a hot
solution of an arsenate in concentrated hydrochloric acid, a mixture of
arsenic trisulfide and pentasulfide is formed.

This behavior is very interesting, but the relations involved are quite com-

plicated. The solubility products of both arsenic trisulfide and arsenic
pentasulfide are extremely small, and it requires but a small quantity of either

230 HYDROGEN SULFIDE GROUP

Ast*+ or Ast++*+ jons to reach this value even with the sulfur ions from
slightly ionized hydrogen sulfide. A cold solution of an arsenate in 0.3- .
normal hydrochloric contains no appreciable quantity of Ast++** cations.
Arsenic acid is of approximately the same strength as phosphoric acid and it
is only in the presence of a very large excess of an acid such as hydrochloric
acid that the ionization of the first hydrogen acid is repressed to a marked
degree. In the presence of concentrated hydrochloric acid, however, it is
reasonable to assume that a small quantity of Ast++** cations are present.
These react with hydrogen sulfide to form the very insoluble pentasulfide,

H;AsO,+ 5Ht o> 4H.O +Ast fk ws 2Ast +++ +4 5HS — As.S;+ 10H*.

The arsenic sulfide is so insoluble that the effect of the acid is, on the whole,
favorable; it favors the formation of Ast+*+** cations and it prevents the
formation of colloidal solutions of As.Ss.

Hydrogen sulfide is absorbed by a cold solution of an arsenate in dilute
acid to a greater extent than can be accounted for by the solubility of hydro-
gen sulfide in water. Soluble thioarsenates are formed:

H;AsO.+ H.S —> H;As0O;8 -- H,0.

Hydrogen sulfide also exerts a reducing effect upon the arsenate. This
reduction takes place very slowly in the cold, but more rapidly if the tempera-
ture of the solution is raised or if the concentration of the H is increased (cf.
p. 228).

H;AsO,+5Ht+S= > Ast+++4H.0+S8; 2Ast+++3H.S > As.S;+6Ht.

As soon as the solution contains an appreciable quantity of either Astt+** or
~Astt* ions, the precipitation of the corresponding sulfide at once takes place.
The temperature of the solution and the concentration of the acid are exceed-
ingly important factors in the precipitation of arsenic by means of hydrogen
sulfide. The pentasulfide is the more insoluble of the two sulfides.

To precipitate the arsenic quickly by hydrogen sulfide from a solution
of an arsenate, without employing considerable hydrochloric acid, it is only
necessary to reduce the arsenic acid by boiling with sulfurous acid, to boil off
the excess of the latter, and then to conduct hydrogen sulfide into the solution,
whereby a precipitate of arsenious sulfide is at once formed.

Arsenic pentasulfide is insoluble in boiling concentrated hydro-
chloric acid, but, like the trisulfide, it is readily soluble in alkalies, ammo-
nium carbonate, and alkali sulfides:

. AsoS;+60H”- =3H20+AsS4 +AsOszS° ;
As2S5+38CO3- =3COz2 T +AsS4= +As038° ;
AsoS;+38S~ =2AsS4 .

By acidifying these solutions, arsenic pentasulfide is reprecipitated:
2AsS4~ +6Ht aoe. 3He2S +AsoSs,
AsS4 +As0O3S +6Ht — 3H20+As2S5.

ARSENIC 231

Arsenic pentasulfide is oxidized by fuming nitric acid to sulfuric
and arsenic acids; also by solution in ammoniacal hydrogen peroxide:

AsoSs5 +20H202+160H =28H20+5804-+ 2AsO04>.

2. Silver Nitrate precipitates from neutral solutions chocolate-
brown silver arsenate (difference from arsenious and phosphoric acids):

AsO4=+3Agt — AgsAsOu,

soluble in acids and in ammonia.

3. Magnesium Chloride precipitates, in the presence of ammonia
and ammonium chloride, a white, crystalline precipitate of magnesium
ammonium arseniate:

AsO4=+ Mgt t4NHat— MgNH.AsOxz.

This precipitate is insoluble in dilute ammonia and is used for the
quantitative determination of arsenic. By ignition it is changed into
Magnesium pyroarseniate:

2MgNH.AsO4 = H20+2NH3+Mg2As207.

4. Ammonium Molybdate, added in considerable excess to a
boiling nitric acid solution, precipitates yellow, crystalline ammonium
arsenomolybdate:

AsO4~ +3NH4t+ 12M00.47+24H*t— 12H2O+ (NH4)3AsO4-12M003.

This precipitate, like that of the corresponding molybdenum com-
pound with phosphoric acid, is insoluble in dilute nitric acid solution
containing ammonium nitrate, but is readily soluble in ammonia or
caustic alkali solutions:

(NH,)3AsO4-12M003+240H~ — 12H20+3NH4*+AsOs- +12M004-,

The yellow precipitate is also soluble in a solution containing an
alkali arsenate; complex anions containing more arsenic, are formed
and the ammonium salts of these complex ions are soluble in nitric
acid. Consequently a large excess of ammonium molybdate should
be used if it is desired to precipitate arsenic acid.

As we shall see later, phosphoric acid behaves similarly toward mag-
nesium salts and ammonium molybdate. If, therefore, phosphoric acid and
arsenic acid are both present, it is necessary to precipitate first the arsenic
with hydrogen sulfide, filter, and oxidize the precipitated arsenic sulfide to
arsenic acid with fuming nitric acid. In such a solution a precipitate produced
by means of ammonium molybdate or magnesium chloride must be caused
by arsenic acid. In the same way a precipitate produced in the filtrate from the
hydrogen sulfide precipitate must be caused by phosphoric acid. This is safer

232 HYDROGEN SULFIDE GROUP

than to depend upon the fact that the ammonium phosphomolybdate forms
more readily at lower temperatures (60°) than does the corresponding arsenic
compound.

5. Potassium Iodide, in a solution strongly acid with hydrochloric
acid, reduces a solution of an arsenate with liberation of iodine (cf. p.
228) ;

HgAsO4+5H*+2I- > Ast +*++4H20-+ Io.

The reaction takes place quantitatively if the iodine is removed —
by adding sodium thiosulphate.

C. Reactions which May be Obtained with All Arsenic Compounds

1. The Marsh Test for Arsenic.—All compounds containing arsenic
may be reduced, in acid solution, by means of nascent hydrogen to
arsine, AsH3:

As203+6H2 =3H20+2AsHs3 | ; As:03+8H2 = 5H20+2AsHs3 T ;
AseS3+6He =3HeS T +2AsH3 t ‘

The sulfides are reduced very slowly, but the oxides are reduced
quickly even at ordinary temperatures. To produce nascent hydro-
gen, zine and sulfuric acid are used.

This very poisonous arsine possesses a property which enables |
us to detect with certainty the merest trace of arsenic—as little as
0.0007 milligram As. By conducting the gas through a heated glass
tube filled with hydrogen, it is decomposed into hydrogen and metallic
arsenic; and the latter is deposited as a brownish-black mirror on
the sides of the glass tube, just beyond the place where it was heated.

This test is extremely sensitive, and must be made with caution,
as almost all reagents, especially commercial zinc and sulfuric acid,
are likely to contain traces of arsenic. Incase these are used with-
out previous testing, arsenic is likely to be found even although it
may not have been present in the substance itself.

The Marsh test is particularly useful for detecting the presence
of very small amounts of arsenic which could not be found by any
of the previously mentioned reactions. In cases of poisoning and
for detecting the presence of arsenic in wall-papers, this test, or a
modification of it, is always used; we will, therefore, discuss it in
detail.

Formation and Properties of Arsine

(a) Formation.—Arseniuretted hydrogen, or arsine, is produced, as above
mentioned, by the reduction of compounds containing arsenic with nascent

ARSENIC 233

hydrogen. For developing the latter, pure zinc and pure sulfuric acid should
be used. If other metals and other acids are used (e.g., tin and hydro-
chloric acid, iron and sulfuric acid), the arsenic compound will be reduced;
but if iron is used, a part of the arsenic is changed to solid arseniuretted hydro-
gen, which remains in the flask and consequently escapes detection. If tin
_ and hydrochloric acid are used, a high temperature is necessary in order to
accomplish the reduction,* while with zine and sulfuric acid the reaction
takes place readily at the ordinary temperature. Chemically-pure zinc
dissolves with difficulty in chemically-pure sulfuric acid, so that it is well to
activate the zinc by the addition of a little foreign metal. The addition of
a drop of chloroplatinic acid causes at first a more rapid evolution of hydrogen,
but the reaction soon slows down and is not accelerated by the addition of
more choloroplatinic acid. Moreover, the addition of chloroplatinic acid
has the disadvantage of causing considerable arsenic to be held back by the
platinum; less than 0.005 mgm. of As.0;, cannot be detected in this way.t
Much better results are obtained by using an alloy of zinc and platinum.
Thus F. Hefti t found that zinc alloyed with 10 per cent platinum caused a
more uniform evolution of hydrogen and that the formation of arsine was
accelerated, while less arsenic was retained by the platinum. With this alloy
quantities of As.O; as small as 0.0005 mgm. can be detected with certainty.
The best activating agent, however, is copper in the form of a zinc-copper
alloy prepared as follows: Melt 20 gms. of the purest zinc in a small Hessian
crucible, stir a very little pure copper into the molten zinc with the aid of a
stick of zinc. Pour the molten metal into water, keeping as much as possible
of the oxide back in the crucible. With this alloy and 15 per cent sulfuric
acid, a steady, continuous current of gas is obtained and it is possible to
detect with certainty as little as 0.00025 mgm. of As.O;.

Arsenic, arsenious oxide, arsenic pentoxide, and arsenic trisulfide are readily
reduced in alkaline solution by sodium amalgam, aluminium, or Devarda’s
alloy and caustic potash, forming arsine. . The reduction takes place quickly,
and the arsine may be detected by the Gutzeit reaction (cf. p. 238). The
presence of organic matter in solutioh hinders the reaction; 3 cc. of urine
in which 1 mgm. of As,O; was dissolved showed no trace of arsine after treat-
ing for hours with Devarda’s alloy and caustic potash solution. In such
cases the organic substance must be decomposed before testing for arsenic.
(Cf. pp. 129 and 152.)

Arsine is also obtained by dissolving many arsenides in hydrochloric
or sulfuric acid:

Zn;As.+6HCl =3ZnCl.+2AsH; T e

The arsenides of iron are attacked by acids only with difficulty, except
when an excess of iron is present, when, with the help of the nascent hydrogen,
they are decomposed, forming solid and gaseous arseniuretted hydrogen.

*Thus VANINO, working at ordinary temperatures, could not detect less than
0.002 gm. of As,O; by means of tin and hydrochloric acid, and where chloroplatinic
acid was added, less than 7; mgm. of As,O; could not be found. Z. angew. Chem.,
1902, 82.

{+ Bernstern, Inaug.-Dissert., Rostock, 1870.

t Inaug.-Dissert., Ziirich, 1907.

234 HYDROGEN SULFIDE GROUP

Consequently iron-sulfide containing arsenic, on treatment with acids, always
yields hydrogen sulfide contaminated with arsine.*

Arsenites can also be reduced to arsine by the action of the electric cur-
rent.’ It is possible to distinguish between an arsenite and an arsenate in this
way.

Certain microbes, namely, Penicilliwm brevicaule, when provided with —
nutriment containing only traces of arsenic, have the power of forming vola-
tile arsenic compound of a garlic odor, and ‘this may be used as an extremely ~
sensitive test for arsenic. By

(b) Properties—Arsine is a_ colorless, unpleasant-smelling, extremely
poisonous gas, which, on being heated away from the air, is decomporrae into
arsenic and hydrogen:

4AsH; =Asy+ 6H, T

By heating in the air, it is oxidized to water and arsenic trioxide. Solid
iodine changes it to arsenious iodide and hydriodic acid:

AsH;+3I, = AsI,;+3HI.

This reaction takes place on conducting arsine over solid iodine. This
property serves to free hydrogen sulfide from arsine, as hydrogen sulfide does not
act upon solid iodine, but only upon aqueous iodine solutions. Arsine is
not attacked by hydrogen sulfide at ordinary temperatures, but at 230° C.
sulfide of arsenic and hydrogen are formed.

Arsine is a strong reducing agent: silver salts are reduced to metal (see
p. 239).

Directions for Performing the Berzelius-Marsh Test

The apparatus devised by G. Lockemann,f shown in Fig. 20, may be used
to advantage. .

In the flask K, of 100 to 150 cc. capacity, place 3 or 4 gms. of zine alloyed
with copper (cf. p. 233) and about 20 ¢c. of 4-normal sulfuric acid free from
arsenic. A steady stream of hydrogen is at once evolved, and in twenty
minutes the air will be entirely driven out of the apparatus. When, at the end
of about twenty minutes, the gas escaping at b is found to be pure (by collect-
ing a little in a small tube and holding it near a flame; it should light without
a sharp explosion), light the hydrogen at b.{ The flame, should be about 2 or 3
mm. high and should remain so during the whole of the experiment; if it
becomes higher, cool the solution in K by placing the flask in cold water,
and, conversely, if the flame is too low add a little more sulfuric acid or place
the flask in warm water. ’

First of all, test the zinc and sulfuric acid to see that they are free from

* (Chem. Zentralbl., 1902, I, p. 1245.)

t Z. angew. Chem., 1905, pp. 427 and 491.

tA safe way to light the flame is to take the small tube used for testing the gas,
fill it with gas, light it, and bring it slowly to the end of the tube b. If the gas is
pure the hydrogen in the small tube will burn quietly for some little time. If
impure, there will be none left in the small tube after it is exploded and this will
not light the escaping gas at 0,

ARSENIC 235.

arsenic. Heat the hard-glass tube at B just before the restriction in the

tube, which is 5 mm. long and 1.5-2 mm. wide. If at the end of twenty minutes
. there is no arsenic mirror formed in this capillary, the ‘Teagents are free from ~

arsenic.

Transfer the sulfuric acid solution to be tested for arsenic, and which must
be free from organic substances,
sulfides, chlorides, nitrates, or other
oxidizing agent, to the graduated
funnel 7’, and add it little by little
to the flask K without in any way
interrupting the current of hydrogen.
Just before adding the solution to
the flask, light the two burners at
A and thereby heat the glass tube
to dull redness. The gas as it
escapes from the flask K_ passes \
through the drying-tube C con-'
taining granular calcium chloride,
and then passes into the tube A,
where any arsine is quantitatively
decomposed into arsenic and hydro-
gen. The arsenic is deposited on
the cold walls of the capillary. Cool
the end of the capillary, in order to
form a sharply defined mirror, by
winding around it a piece of wicking, as shown in Fig. 20, and allowing
water to drop upon it from the dish W during the experiment.

All the arsenic will be deposited at the end of an hour, and by comparing
the mirror with a series of standards the amount can be estimated accurately
(see page 238).

Remark.—If{ the tube A is
not heated at all, but the gas
ignited at b as above described,
the arsenic may be deposited
upon a cold porcelain dish by
holding the dish in the flame.
The deposit is readily soluble
- in sodium hypochlorite solution
(difference from antimony). In
this form the test was used by
James Marsh in 1836.

Confirmatory Test.—In_ the

Fic. 21. small glass tube open at both

ends (see Fig. 21) the arsenic

mirror is found. Hold the tube in an inclined position and heat it by

means of a small flame whereby the arsenic is changed to arsenic trioxide,

giving off the characteristic garlic odor, which can be detected at the upper

end of the tube if only ;4;5 of a milligram of arsenic trioxide is formed.

After the tube is cooled, the arsenic trioxide is to be found at a in the form

of small glistening octahedrons, which can be seen with the magnifying glass
or often with the naked eye.

236 HYDROGEN SULFIDE GROUP

These three facts—formation of the mirror, the garlic odor, and the
octahedrons—suffice to prove the presence of arsenic; but the more proofs we
have, the more certain we are of the accuracy of the result. If the octa-
hedrons have been recognized, seal the capillary end of the tube with a flame,
and introduce 1 to 2 drops of pure, con-
centrated hydrochloric acid into the tube
with the help of a dropper, and move the
tube so that the arsenic trioxide is moist-
ened by the acid; then add 6 to 10 drops
of distilled water and pass hydrogen sul-
fide into the tube, whereby yellow arseni-
ous sulfide is formed.

The hydrogen sulfide required may be
generated from a solution of sodium sulfide
by allowing it to flow into dilute sulfuric
| acid, as illustrated in Fig. 22. The upper
¥ N4cotton Part of the test-tube contains a wad of
cotton wool, which prevents any of the
solution in the tube from being mechani-
cally carried over into the tube containing
the arsenic.

As an example of the practical appli-
cation of this delicate test, we shall de-
scribe the method to be employed in the
detection of arsenic in wall-papers, etc.
The amount of arsenic contained in wall-
papers is usually so small that weighing
the mirror produced would not be accu-
rate.* It is best, therefore, to prepare a
number of mirrors from known amounts
of arsenic, to establish a scale for deter-
mining how much is contained in the given wall-paper or fabric.

First of all, the arsenic must be extracted completely from the paper, and
to this end it is necessary to decompose the organic material, which is accom-
plished preferably as follows:

Na,$

Decomposition of Organic Material

Take exactly one square decimeter (100 sq.cm.) of wall paper, roll it into
a cylinder and push it down into a tube closed at one end, such as is used for
the Carius determination of the halogens (see Volume II of this book). Add
2 cc. of pure, fuming sulfuric acid (25 per cent oleum, Kahlbaum) through a
long-stemmed funnel. Then pour 3 or 4 cc. of fuming nitric acid into a small
test-tube and carefully allow the latter to slip down the sides of the Carius
tube so that the two acids-do not come in contact with one another. At
the open end of the tube draw out the glass to form a strong capillary and

*In Massachusetts, the law permits the presence of 0.1 mgm. per square meter
in wall-paper, but only 0.01 mgm. per square meter in wearing apparel. In most
cases it is merely necessary to determine whether the legal limit is exceeded.

7 C. R. Sanarer, Amer. Acad. of Arts and Sciences, 26, 24.

ARSENIC 237

seal the end as described in Volume II. Cover the tube with asbestos paper
and slowly heat it inside a strong iron tube, in the furnace used for Carius
tubes, to a temperature of about 230° and keep it at this temperature for
an hour and a half. Allow the tube to cool and then withdraw it from the
iron protective tube, by means of a wire previously fastened to it, until the
capillary projects a little; heat this with a Bunsen flame. As soon as the glass
becomes soft the pressure on the inside of the tube blows out a hole through
which gas escapes until the pressure is the same inside the tube as without.
Break off the point of the tube and rinse the contents of the tube and of the
tip into a porcelain evaporating dish. The colorless solution * thus obtained
will contain all the arsenic in the form of arsenic acid. Evaporate the solu-
tion till fumes of sulfuric acid are evolved thickly, and then, after cooling,
add 15 cc. of water and pour the liquid into the funnel 7 of Fig. 20, rinsing
the dish twice with 3 cc. portions of water. After mixing the liquid in 7 by
means of a small stirring rod, make a note of the total volume of the
liquid.| During this operation, the Marsh apparatus should be made ready
for the test.

When the apparatus is ready, add a few drops of the well-mixed solution
through 7 to the reduction flask K. If no mirror appears within three or
four minutes, add one-eighth to one-quarter of the filtrate little by little;
and if no mirror appears after five minutes, the whole filtrate. The whole
filtrate is not added at once, because if too strong a mirror is obtained, it is
much more difficult to estimate the amount of arsenic present. After twenty-
five minutes all the arsenic will be deposited if not more than 0.05 mg. of
arsenic is in solution. If a mirror of sufficient density was cbtained in fifteen
minutes from only a fraction of the whole solution, no more should be added,
but the operation should be continued for ten minutes more, the flame
extinguished, and the tube allowed to cool while hydrogen continues to pass
through it. The mirror is then compared with the scale, and the remaining
part of the filtrate is weighed in order to determine how much was used for
the test.

If sufficient material is at hand, a duplicate experiment should be made
with a new tube and a new sample. The results of a few such determinations
are given in the following table:

Total Weight

2 i i :

Sees | eet || eee | ee) eee
100 31.63 31.63 0 0 0
100 30.11 10.23 0.015 0.044 4.4
arene cla) ree. ne's 9:87 0.013 0.0399 3.99
100 28.72 8.32 0.045 0.155 15.5
Choate SEE eck sales 7.53 0.042 0.163 16.3

50 30.22 2.64 0.015 0.172 34.4
ree REMC ey 3.22 0.020 0.187 37.4

*If the paper contained iron or aluminium compounds, the anhydrous sulfates
are formed. These dissolve by heating with water. The above method is the
best and cleanest for decomposing organic material.

+ Instead of measuring the liquid it may be weighed.

238 HYDROGEN SULFIDE GROUP

The comparison of the mirrors is best made in transmitted light. The
normal mirrors are prepared as follows: Dissolve 1 gm. of pure, sublimed
arsenic trioxide in a little sodium carbonate solution, acidify with dilute sul-
furic acid, and dilute to a liter. Take 10 ce. of this solution, of which ‘1 ee.
contains 1 mg. of As,O;, and again dilute to a liter; in this way a solution
is obtained of which 1 cc. contains exactly 0.01 mg. of As.O;. Measure out
1 cc., 2 ec., 3 cc., 4 ec., and 5 cc. of the solution and introduce each portion
separately into the Marsh apparatus, and prepare the corresponding mirrors
in different tubes. It is best to prepare two tubes from each amount of arsenic,
as the mirrors are not always the same. These mirrors may be kept in the dark
for some time; but on exposure to the light they fade perceptibly. Mirrors
which are sealed up with hydrogen do not keep as well.

Testing Urine, Blood, Milk, Beer, etc., for Arsenic

Evaporate 100 cc. of the liquid in question to dryness in a porcelain dish.
By means of a spatula introduce the residue as completely as possible into a
tube, such as is used for the Carius determination (Vol. II), and add 4 ee.
of a 25 per cent oleum. In order to transfer the rest of the residue to the
tube, pour 2 cc. of fuming nitric acid into the dish and after wetting all the
sides of the dish, transfer the acid to’ a small test-tube. Repeat this operation —
twice more and then allow the small test-tube to slip into the Carius tube.
Draw out a capillary at the open end of the tube and seal it by the flame.
Heat the tube in the Carius furnace for one hour at 160°. After cooling, —
open the point of the tube with the usual precautions (p. 237), and release
the pressure. Seal the tube again and heat for half an hour to an hour at
230°. The operation is then continued as described above.

With urine a somewhat different procedure is followed. Evaporate liquid
not quite to dryness, but to sirupy consistency; spread the sirup upon a porce-
lain boat and allow it to slip into the Carius tube. The rest of the process is
carried out as before.

2. The Gutzeit Test for Arsenic depends upon the behavior of
arsine toward a concentrated solution of silver nitrate (1 : 1) (accord-
ing to Eidenbenz, a crystal of solid silver nitrate should be used).
The silver nitrate is at first colored yellow and then black, the follow-
ing reactions taking place: :

1. 6AgNO3+AsH3 = AsAgs -8AgNO3+3HNOs.
Yellow

2. AsAg;-3AgNO3+3HOH = H3As03+3HNO3+6Ag.

The test is carried out as follows: Place a little of the substance in a
small test-tube, Fig. 23, add a few grains of zinc and a little dilute sulfuric
acid, and place a wad of cotton near the top of the tube asa filter. Over the
mouth of the tube place a piece of filter-paper with a crystal of silver nitrate
on top.

If arsenic is present, the silver nitrate is at first turned yellow, but it
becomes black very quickly.

ARSENIC 239

This reaction is often used for quickly testing commercial acid for arsenic,
but it is not as reliable as the Bettendorff test (p. 228), because phosphine *
and stibine give a similar reaction with silver nitrate,
while they are not reduced by stannous chloride. AgNO,

If arsine is allowed to act upon a dilute solution of Fans \ Paper
silver nitrate, the yellow compound AsAg;-3AgNO; is
not formed, for it is immediately decomposed hydro- oh

@| Cotton

lytically, according to the equation BUEN
(aaa) wD

AsH,+6AgNO;+3HOH =6HNO,+H,AsO;+6Ag. teh

If the precipitated silver is filtered off, and am-
monia then poured on top of the filtrate, the neutral
zone will appear yellow owing to ae formation of
silver arsenite.

This reaction never takes place ‘iit quantita-
_ tively; the deposited silver invariably contains a little
silver arsenide, Ag;As.

If the silver solution is made ammoniacal, it is
true that all the arsine .will be absorbed, but the de-
posited silver still contains a little arsenic and the
solution a small quantity of ammonium arsenate. If, Fig. 23.
however, the ammoniacal solution containing the
silver nitrate and ammonium. arsenite is heated to boiling, then the arsenite
is oxidized quantitatively to arsenate with deposition of silver.

2Agt+AsO;-+OH~ — HAsO,-+ 2Ag.

Under these conditions eight atoms of silver are deposited from each origi-
nal molecule of arsine:

AsH;+8Agt+110H~ — AsO,=+7H,0-+ 8Ag.

If the deposited silver is filtered off and the filtrate carefully neutralized
with nitric acid, a brown precipitate of silver arsenate is formed.

Somewhat less sensitive than the original Gutzeit test, although
very satisfactory, is the modification recommended by Fliickiger ft
and Lehmann.§

Instead of allowing the arsine to act upon silver nitrate, bring it into contact
with mercuric chloride || paper, which is turned yellow by a little arsine and
reddish brown by considerable arsine. {]

* Commercial zine often contains a small quantity of phosphorus.

{ Cf. REcKLEBEN, LocKEMANN and Eckuarpt, Z. anal. Chem, 1907, 671.

t Arch. Pharm. (31, 1889, 27.

_ § Pharm. Zig. Berlin, 1892, 36.

|| To prepare the mercuric chloride paper moisten some filter paper with an
alcoholic solution of mercuric chloride, allow the alcohol to evaporate and repeat
the process four or five times.

q Aueust Gorruetr, J. Soc. Chem. Ind., 22, 191 (1903).

240 HYDROGEN SULFIDE GROUP

The exact composition of these compounds is not known definitely.
Possibly As(HgCl); is formed first; and then, by further action of AsHs,
AsH(HgCl)2 and As.Hg; are formed.*

AsH3+3HgClp— 3HCl+ As(HgCl)3
2As(HgCl)3+ AsH3— 3AsH(HgCl)>
As(HgCl)3+ AsH3 — 3HCl+ AsoHg3.

These arsenic compounds are characterized by their insolubility in 80 per
cent alcohol.

Stibine gives no reaction in this test when little of it is present, but the
presence of somewhat more of it causes the formation of a brown spot which
is soluble in alcohol. If, however, arsenic and antimony are both present,
the former is recognized by cutting out the spot from the rest of the filter-
paper and placing it in 80 per cent alcohol, whereby the brown spot due to
the antimony is removed in a short time and the yellow arsenic spot appears
plainly. When considerable antimony is present, the test fails; a gray-black
spot is produced which does not disappear on treatment with alcohol.

3. The Reinsch Test is very easy to make, but it is not as sensitive
as the tests just mentioned. It depends upon the fact that when
a strip of polished copper foil is added to a solution of arsenious acid,
the copper is colored gray owing to the deposition of CusAse on the
copper. |

From concentrated solutions the arsenic separates out in the
cold, but from dilute solutions only on warming. If considerable
arsenic is present, the gray copper arsenide drops off from the copper.
Antimony is also precipitated on copper from its solutions, so that
the deposit must be tested for arsenic in the dry way. Arsenic acid
is also reduced by copper, but. only on warming.

The Reinsch test is often used in testing wall-papers for arsenic.
The pieces of paper are treated with a little hydrochloric acid (1 : 2),
a piece of copper foil added, and warmed. A gray deposit on the copper
indicates the presence of arsenic.

To confirm this test, the piece of gray copper foil is placed in a
tube of difficultly fusible glass and heated in a stream of hydrogen
gas; an arsenic mirror is produced which can be tested as described
on p. 235. |

Detection of Arsenic in the Human Organs

To detect small quantities of arsenic present in the organs of persons who
have probably died from poisoning, it is necessary in the first place to destroy

* Besides the above compounds, AsH2(HgCl) and As,Hg; are said to exist.
Cf. Parrueit, Arch. Pharm., 237, 121.

ARSENIC . 241

all the organic tissue; this may be accomplished by the action of sulfuric and
nitric acids.* Place 200 g. of the organ, 200 cc. of concentrated nitric acid and
5 ec. of 2 per cent permanganate solution in a 2-liter flask and heat until
foaming ceases. Then transfer tlhe solution to a 1-liter flask, rinsing out
the larger flask with 100 cc. of concentrated nitric acid and 100 cc. of water.
Boil four hours, or until the solution is reduced to 80 cc. in volume. Add 100 ce.
of concentrated sulfuric acid and evaporate till fumes of sulfuric acid are -

evolved. Cool, add 2 or 3 cc. of concentrated nitric acid and again heat until

white fumes are evolved and repeat this treatment with nitric acid about six
times. Then, when all the nitrie acid has been expelled, cool, dilute with
100 cc. of water and test the solution by the Marsh test.

To determine the quantity of arsenic present, transfer the solution to a dis-
tilling flask, add 20 ce. of concentrated sulfuric acid, 50 gms. of sodium chloride,
1 gm. of potassium chloride and 10 gms. of ferrous sulfate crystals. Distill
into a flask containing 50 gms. of sodium bicarbonate and 100 cc. of water
until all the solid bicarbonate dissolves. Titrate the sodium arsenite solution
thus obtained with iodine as described in Vol. II. A blank experiment should
be made with all the reagents to make sure that they are free from arsenic.

REACTIONS IN THE DRY WAY

Metallic arsenic burns, giving off a garlic odor. Mixed with sodium

carbonate and heated on charcoal, all arsenic compounds give this

odor.

Oxygen compounds of arsenic are easily reduced to metal in the
upper reducing flame. On holding a porcelain dish (glazed on the
outside and filled with water) directly over the sample, the arsenic
vapors are condensed on the dish, forming a brownish-black coat-
ing which is soluble in sodium hypochlorite solution, disappearing
instantly, the arsenic being oxidized to arsenic acid:

Aso+5Na0Cl+3H20 =5NaCl+2H3AsOg.

If the porcelain dish is not helé closely above the reducing flame,
but above the upper oxidizing flame, the arsenic vapors are burned
with a bluish flame to white arsenious oxide which deposits on the
dish.

If this deposit is moistened with silver nitrate, and ammonia
vapors blown upon it, a yellow ccloration due to AgsAsOz is formed,
which disappears if more ammonia is allowed to act upon it (differ-
ence from antimony):

* Cf. Gautier, Bull. soc. chim. Paris, 25, 252 (1875); CurrreNDEN and DoNaLp-
son, Am. Chem. J., 11, 236 (1880-1); JoacnimoGuti, Arch. exp. Path. Pharm., 78,
1-16 (1914).

242 HYDROGEN SULFIDE GROUP
AseO3 + 6AgNOs3 + 3H2O = 2Ag3AsOz3 + 6HNOs.

The ammonia serves to neutralize the nitric acid formed by the:
reaction, but the precipitate dissolves in excess of ammonia as well
as in nitric acid.

The microchemical method of Hartwich and Toggenburg* is often useful
when the arsenic is present as trioxide. Prepare a glass cylinder about 12 mm.
in diameter and 10 mm. in height and make sure that both the upper and
lower edges are smooth. Place this cylinder upon a small watch-glass and
pour into it a little of the substance to be tested, well mixed with ignited sand.
Cover the cylinder with a flat glass slide about 50 mm. square. Heat the
watch-glass very carefully with the flame from a small burner; the flame
should not be over 5mm. high and 30 to 40 mm. below the watch-glass. After
heating ten or fifteen minutes, allow to cool slowly and finally examine the
bottom surface of the glass slide. If arsenic is present as trioxide, it will have
sublimed and the vapors will have condensed, upon the slide for the most part,
in the form of beautiful octahedrons, the shape of which is very distinct when
viewed under the microscope. These crystals are visible when only 0.01 mg.
of arsenic is present. The sublimate may be identified further by the silver
reaction. |

ANTIMONY, Sb. At. Wt. 120.2
Sp. Gr. =6.7-6.8. M. Pt.=630°C. B. Pt.=about 1450° C.

Occurrence.—Antimony seldom occurs free in nature, although
large amounts of the metal have been found recently in Australia.
The most important compounds containing antimony are (as with
arsenic) the sulfur compounds. Stibnite, SbeS3, orthorhombic, is
found in Japan in beautiful crystals. The occurrence of kerme-
site, Sb2OSe, is interesting, as this compound is often met with in
analysis.

Of the oxygen compounds the dimorphous antimony trioxide is
known as isometric senarmontite and orthorhombic valentinite.
Antimony also occurs in many thio salts, of which the tribasic silver
thioantimonite, or pyrargyrite, AgsSbSs3, may be mentioned.

Antimony is a silver-white, brittle metal. It burns readily in
the air to antimony trioxide. The solvent for antimony is aqua regia,
by which it is converted into chloride. Nitric acid attacks antimony,
changing it into Sb2O4, which dissolves slightly in concentrated acid,
but is insoluble in dilute acid.

* Schweiz. Wochenschrift fiir Chem. u. Pharm. 1909, No. 52, p. 1.

ANTIMONY 243

Antimony forms three oxides: antimony trioxide, Sb2O3; anti-
mony pentoxide, Sb205; and antimony tetroxide, Sb2O04, which may
be regarded as antimonous antimonate, and is a very indifferent
substance chemically. Antimony trioxide as a rule shows basic prop-
erties, while antimony pentoxide has more the character of an acid
anhydride.

A. Compounds of Antimony Trioxide

By burning the metal in the air, the trioxide is obtained, which
on stronger ignition in the presence of air is changed to the inert
Sb20..

The trioxide is dissolved by concentrated hydrochloric acid, form-
- ing antimony trichloride, a compound which (like bismuth chloride)
is readily changed into a basic salt by the action of water, the decom-
position of which depends upon the concentration of the reacting sub-
stances. Thus an oxychloride SbOCl is known, which is formed
according to the following equation:

SbCl3 + H20 = 2HCl+SbOCI1.

In the presence of a large amount of water some oxide is also formed
with the oxychloride:

2SbOCI+H20 = 2HCI1+Sb20s.

A mixture of oxychloride and oxide is known as “ algarot ”’ powder,
Sb203-2SbOCI.

By boiling with considerable water the oxide alone is obtained.

Antimony trioxide forms three hydroxides, which behave as very
weak acids: orthoantimonous acid, H3SbOs; pyroantimonous acid,
H4Sb205; and the hypothetical metantimonous acid, HSbOs.

Salts of the metantimonous acid are known, although the free
acid itself has never been isolated. On boiling the oxide Sb2O3 with
concentrated caustic soda or potash, it goes into solution, but on
dilution with considerable hot water Sb203 separates out again. On
filtering this off, tetragonal crystals of NaSbOeg are deposited in the
filtrate on cooling; they are, however, very unstable, and are decom-
posed by standing in the air into sodium carbonate and antimony
trioxide. By dissolving antimony trioxide in strong alkali, the ortho-
antimonate is probably formed,

Sb203+60H- — 2SbO3 +3H20,

244 HYDROGEN SULFIDE GROUP

which is hydrolyzed on dilution into metantimonite and alkali hy-
droxide:

SbO3=+H20 — 20H-+S8b0.7.

The latter is decomposed by more water into trioxide and alkal:
hydroxide; so that on adding to a solution of the trichloride either
sodium hydroxide or carbonate, an almost quantitative precipita-
tion of Sb2Oz will be obtained:

2SbCl3 +6NaOH = 6NaCl+3H20+Sb203;

2SbCl3+3Na2CO3 =6NaCl+3CO02+Sb203.

Antimony oxychloride, SbOCI, contains the univalent group, SbOT,
which is known as the antimony] group.

Antimony oxychloride, therefore, can be regarded as antimonyl
chloride. Antimonyl nitrate SbO(NOs), is also known, and anti-
monyl sulfate, (SbO)2SO4. All these compounds are easily hydrolyzed
into acid and oxide, so that they are rarely met with in the course of
analysis, with the exception of antimony] chloride. )

The antimonyl compounds of certain organic acids (such as tar-
taric acid) are very much more stable. :

On boiling antimony trioxide with a solution of potassium acid
tartrate, it goes readily into solution, forming the so-called ‘ tartar
emetic,”

2KHC4H406+Sb203 — H20+2K(SbO)C4H40¢,

which is comparatively soluble in water. 100 cc. of water dissolve:
5.3 gms. at 8.7°; 7.9 gms. at 21°; 12.2 gms. at 31°; 18.2 gms. at 50°,
and 31.2 gms. at 75°.
Not only antimonyl oxide, Sb2O3, but all antimonyl compounds,
form a complex anion with tartaric acid or with a tartrate, thus,
antimony] chloride dissolves in Rochelle salt, or in tartaric acid:

C4H40¢6~+SbOC1 — Cl +[(SbO)CaH40¢)-;
H2C4H406+SbOC1 — HCI+ H[(SbO)C4H 40g].

Tartar emetic, K(SbO)C,H.,0.-3H.O, is the most important antimony
compound of commerce. Consequently it will be worth while to say a few
words with regard to its behavior toward acids.

ANTIMONY | 245

If an aqueous solution of potassium antimony] tartrate is treated with
hydrochloric acid, a white precipitate of antimony] chloride ,is formed,

K(SbO)C,H,0.+2H* — H.C,H,0,+ K++ SbOCI,
which readily dissolves in more hydrochloric acid,
SbOCI+ 2HCl = H:0+S8bCl,,

but, on the addition of more water, it is reprecipitated, etc.

Sulfuric and nitric acids precipitate orthoantimonous acid from a solution
of potassium antimony] tartrate; for the antimonyl compound, which is at
first formed, is immediately decomposed by water,

[(SbO)C,H,0.]-+Ht+2H.0 << H.C,H,O,+ H;SbO;.

REACTIONS OF ANTIMONOUS COMPOUNDS IN THE WET WAY

‘1. Water precipitates at first a basic salt which is changed into
oxide by more water.

2. Sodium Hydroxide, Ammonia, and Alkali Carbonates precipi-
tate the amorphous hydrated oxide.

3. Hydrogen Sulfide precipitates, from solutions which are not
too acid, flocculent, orange-red antimony trisulfide:

2Sb++*+3H2S — 6H*+SbeS3.

As is indicated in the equation, the antimony trisulfide is in equilibrium with
hydrogen ions; in 12-normal acid it is readily soluble (difference from arsenic).
On the other hand, antimony sulfide is less soluble in acid than is tin sulfide
and it can be precipitated from a solution which is normal in acid concentration.
The trisulfides of both tin and antimony are somewhat more soluble than the
pentasulfides. If a solution of antimony in concentrated hydrochloride is to
be treated with hydrogen sulfide and the solution is first diluted, a precipitate
of SbOCL is likely to form. This does no harm, as hydrogen sulfide will change
the oxychloride into the less soluble trisulfide. If it is desired to precipitate
the sulfide without any formation of oxychloride, the solution should be satu-
rated with hydrogen sulfide before it is diluted, then diluted and again satu-
rated with hydrogen sulfide.

Antimony trisulfide is soluble in ammonium sulfide, forming a
soluble thio salt:
Sb2S3+3S7 — 2[SbS3]~.

The triammonium salt has never been isolated, the mono salt, NH,SbS,,
alone being known in the solid state. In solution, however, particularly in
the presence of considerable ammonium sulfide, the SbS;= ion is probably
present.

246 HYDROGEN SULFIDE GROUP ,

If yellow ammonium sulfide is employed, ammonium salts of
thioantimonic acid are obtained:

SboS3+2(NH4)2S2 = (NH4)3SbS4-+ N H4SbSs3.

If the solution of ammonium thioantimonite is boiled for a long time
in the air, the red-colored oxysulfide is often precipitated:

2(NH4)3SbS3+402 = 2(NH4)28203 +2NH3+H20+Sb2820.

By boiling antimony chloride with sodium thiosulfate, the oxy-
sulfide is also obtained,

2Sb*t *+*++358203" > 4802 7 +Sb2820,

which, on being warmed with ammonium sulfide, redissolves, forming
the thio salt. |

Antimony trisulfide is also soluble in caustic alkali, forming thio
and oxythio salts:

SbeS3+20H™ = H20+O8bS+S8S8bS-.

These thio salts are decomposed by acids, precipitating antimony |
trisulfide, with evolution of hydrogen sulfide, b

2SbS3 +6H* — 3H2S Tf +Sb2Ss;
2SbS4=-+6H* — 3H28 t +SbeSs;
OSbS~+SSbS~+2H* — H20-+Sb2Ss.

4. Zinc precipitates from solutions of antimony compounds metallic
antimony. If a piece of platinum foil and a little zine are placed in
an antimony solution containing hydrochloric acid, so that the two
metals touch one another, the antimony is deposited on the platinum
in the form of a black stain which does not disappear on removal of
the zine (difference from tin).

Lead will also cause the same reaction to take place (cf. p. 44).

5. Potassium Iodide does not set free iodine when treated with
an antimonous solution (difference from antimonic compounds).

B. Antimonic Compounds

Antimony pentoxide, Sb2O5, is formed as a yellow powder by oxi-
dizing antimony with concentrated nitric acid and gently igniting
the reaction product (antimonic acid). On strong ignition it loses

oxygen and goes over into the very stable antimonous antimonate
Sb204.

ANTIMONY 247

The pentoxide dissolves in concentrated hydrochloric acid, form-
ing the pentachloride

Sb20;+10HCl=5H20+2SbCls.

If this solution is treated with water, a white precipitate of anti-
monic oxychloride, SbO2Cl, is formed, which, by the addition of more
water, is changed on warming into antimonic acid:

SbCl; -+2H20 = 4HC1+Sb0.Cl,
Sb0.Cl+2HOH @ HC1+H;Sb0u.

Tartaric acid prevents the precipitation of the oxychloride, as
with SbOCI (p. 244). Antimony pentoxide is an acid anhydride,
and, like the corresponding P205, can be referred to three acids’
orthoantimonic acid, HsSbO4; metantimonic acid, HSbO3; and
pyroantimonic acid, H4aSbeO7, which have all been isolated. The
salts of the metantimonic and pyroantimonic acids are the most com-
mon. ‘The trimetallic salts of the ortho acid have never been isolated,
but the monometallic salts are known to exist. All antimonates,
being salts of a weak acid, are very unstable, being easily hydrolyzed
by water. ¢

If antimony pentoxide is fused with an excess of caustic potash,
the product of the fusion probably contains the trimetallic salt of
orthoantimonic acid. If, however, the melt is dissolved in a little
water and allowed to crystallize, deliquescent crystals of potassium
pyroantimonate, K4Sb207, are formed.

The ortho salt, which is first formed, is decomposed by water as
follows:

2K3SbO4+ H20 @ 2KOH+ KaSbe207.

By the action of considerable cold water (or more quickly by
rapid boiling with less water) this deliquescent salt is gradually
changed into the acid salt, losing KOH,

K4Sbe07+2H20 @ 2KOH+Ke2H2Sb20;7,

which separates out with 6 molecules of water as a granular powder,
difficultly soluble in cold water. It dissolves to a considerable extent
in water at about 40°—50° C.; this solution is used in testing for sodium,
as the sodium salt is very much more insoluble in water.

On boiling the granular potassium salt for a long time with con-
siderable water, it gradually takes on water, forming the easily soluble
monometallic salt of orthoantimonic acid,

K2H2Sb207+H20 @ 2KH2SbO04,

2A8 HYDROGEN SULFIDE GROUP

which is obtained, on evaporating the solution, as a gummy mass
of the composition 2KH2SbO4+H20, but on boiling the aqueous
solution for a long time, more KOH is lost, with the formation of
amorphous orthoantimonic acid: |

KH2SbO4+ H20 @ KOH+H3SbO..

All antimonates are decomposed by acids, amorphous antimoniec
acid separating out.

The gummy, monometallic salts give an amorphous precipita-
tion with sodium salts, gradually becoming crystalline, while the
potassium pyroantimonate gives a crystalline precipitate immediately.

The relations of antimonous to antimonic salts and of antimonites to anti-
monates are, on the whole, similar to the corresponding relation with arsenic.
In the higher state of oxidation the acid properties are more pronounced than

the basic properties, but antimony, being lower in same group of the periodic
table, is more basic than arsenic.

REACTIONS OF ANTIMONIC COMPOUNDS IN THE WET WAY

1. Hydrogen Sulfide precipitates from fairly acid ‘solutions the

orange-red pentasulfide:
+44

28b 4 +5HeS a 10Ht+SboSs.

Antimony pentasulfide is soluble in 12-normal hydrochloric acid,
forming antimony trichloride, with deposition of sulfur and evolution
of hydrogen sulfide:

Sb2S5;+6HCl = 3H2eS+S8e2 +28bCls.
It also dissolves (like the trisulfide) in alkali sulfides, and in alka-

lies, but not in ammonium carbonate. By treatment with an alkali
sulfide, a thio salt is obtained,

Sb2S5+38~ — 28bS.4°,

which is decomposed by the addition of acids, forming the insoluble
pentasulfide with evolution of hydrogen sulfide:

2SbS4-+6H* — 3H2S fT +Sb2Ss.
Alkalies dissolve the pentasulfide, forming thio and oxy-thio salts:
Sb2S5+60H~ — Sb84=+8-SbO3 +3H20.

2. Hydriodic Acid reduces antimonic compounds in acid solu-
tions, with separation of iodine (difference from antimonous com-
pounds) :

Shi ttt t+91- + Spttt+tp.

a a | a

ae og
Pd!

TIN 249

_ 3. Nascent Hydrogen.—By treating any solution which ¢éontains
antimony with nascent hydrogen, stibine is formed:

Sb203+6H2 =3H20+2SbHs f ,
Sb205-+8H2 = 5H20+2SbHs } .

If the stibine is generated in a Marsh apparatus (cf. p. 235), and the gas
is conducted through a red-hot glass tube, a mirror of metallic antimony will
be deposited, as with arsenic. But as stibine is much more unstable than
arsine, and the antimony itself is much less volatile, the mirror is found nearer
the heated place than is the case with arsenic—sometimes before the hottest
part of the tube is reached—as the decomposition of the stibine takes place
at a much lower temperature than with arsine.

If the stibine is allowed to escape from the tube with the hydrogen, it
' burns with a pale greenish-white flame to water and antimony trioxide. If
a piece of glazed porcelain is held directly over the flame, a deposition of
metallic antimony is obtained which is unaffected by a solution of sodium
hypochlorite (difference from arsenic).

If stibine is allowed to act upon a solution of silver nitrate, a black pre-
cipitation of silver antimonide is thrown down:

SbH;+3Agt =Ag;Sb+3H".

Solid silver nitrate is turned yellow at first, then black; exactly the same
as by arsine (cf. p. 238).

REACTIONS OF ANTIMONY IN THE DRY WAY

Antimony compounds impart to the flame a pale, greenish-white
color. Heated with sodium carbonate on charcoal, a brittle metal-
lic button is obtained, surrounded by a white incrustation.

Compounds containing oxygen are reduced in the upper reducing
flame to metal, which is volatile and burns in the upper oxidizing
flame to trioxide; the latter can be deposited on a glazed porcelain
surface. If the deposit is moistened with silver nitrate solution,
and ammonia blown upon a it becomes black, owing to the separation
of metallic silver:

Sb203-+4AgNO03+4NH3+2H20 = 4NH4NO3 +Sbe20; +4Ag.

TIN, Sn. At. Wt. 118.7
' Sp. Gr. =7.29. M. Pt. =232°C. B. Pt.=about 1500° C.

Occurrence.—Tin does not occur free in nature, but mostly in
‘the form of the dioxide, as tetragonal tinstone, or cassiterite, isomor-
phous with rutile (TiOg), zircon, and polianite (MnOz).

Tin is a silver-white metal, which is ductile and malleable at ordi-

250 HYDROGEN SULFIDE GROUP ~

nary temperatures, but at low temperatures and near the melting-
point it is so brittle that it can be powdered. In order to pulverize
tin, heat it in a porcelain dish till it melts, remove- the flame and
quickly crush the substance with a pestle. It soon cools to about
200°, becomes brittle and yields a fine powder.
Tin is soluble in hot, concentrated hydrochloric acid with evolu-
tion of hydrogen:
Sn+2H*t — Sntt+He.

In the presence of platinum the solution takes place more quickly

and at a lower temperature. Dilute hydrochloric acid dissolves tin, -

but very slowly.
Nitric acid, of sp.gr. 1.2 to 1.3, does not dissolve tin, but oxidizes
it to metastannic acid:

3Sn+4HNO3+H20 =3H2Sn03+4NO f.

Cold, dilute nitric acid dissolves the metal very slowly, without any

evolution of gas, forming ammonium and stannous nitrates.

In this reaction the tin is given two positive charges and a part of the
nitrogen of nitric acid is reduced from a positive valence of five (toward oxy-
gen) to a negative valence of three (toward hydrogen) thereby losing eight
charges. Thus one molecule of nitric acid oxidizes four atoms of tin and
eight more molecules of nitric acid are required to form stannous nitrate and
one to form ammonium nitrate. The whole reaction may be expressed as
follows:

4Sn-+ 10HNO3 = 4Sn(NOs) 2 +NH4NO3 +3H20.
Aqua regia dissolves tin, forming stannic chloride:
3Sn+4HNO3+12HCl1=4NO 7 +8H20+3S8nCh.

Tin dissolves in dilute sulfuric acid very slowly, but readily in
hot concentrated acid, forming stannic sulfate, with evolution of
sulfur dioxide: |

Sn+4H2SO04= 2802 fT +4H20+Sn(S04)2.

Tin forms two oxides: stannous oxide, SnO; and stannic oxide,
SnOz. Salts are known corresponding to both these oxides—stannous
and stannic salts. The former contain bivalent tin and the latter quad-

rivalent tin. Tin is more basic in its properties than antimony, —

but stannous hydroxide is slightly amphoteric (cf. p. 52) and the
acid properties of stannic hydroxide are more pronounced.

| —— _

TIN 251

Stannous Compounds

Stannous oxide (according to the way it is prepared) is either an
olive-green or a black powder, which, on being warmed in the air,
like all stannous compounds readily changes to stannic oxide. By
dissolving stannous oxide (or, better still, the metal itself) in hydro-
chloric acid, stannous chloride is obtained, which is the most im-
portant of all the stannous salts. This salt, with two molecules of
water of crystallization, SnCle+2H20, is the so-called “ tin salt” of
commerce.

Fresh crystals of ‘ tin salt ”’ will dissolve clear in a little water;
if more water is added the solution becomes turbid, owing to the
formation of a basic salt,

SnClo+H20 2 HCl+Sn(OH)CI,

which is readily soluble in hydrochloric acid.
The clear concentrated solution also becomes turbid on stand-
ing in the air, owing to the formation of the same basic salt.

28nClo-+0-+H20 > 2Sn(OH)CI1+Cle,

The chlorine, however, is not set free, but unites with some of the
stannous ions, forming stannic ions:

Sn**--Clh —s Sat t?+-LOCr-.

If tin tetrachloride is treated with metallic tin, the latter goes
into solution and the former is reduced to stannous chloride:

SnCl4+Sn = 28nCle.

Consequently, in order to keep a solution of stannous chloride
in the stannous condition, hydrochloric acid should be added to pre-
vent the formation of the basic salt, and metallic tin to keep the solu-
tion reduced.

Such a solution constantly grows more concentrated, owing to
the gradual solution of the tin. In order to keep a solution of stannous
chloride at a definite concentration (only necessary for purposes of
quantitative analysis) the hydrochloric acid solution should be kept
out of contact with air in an atmosphere of carbon dioxide without
the addition of metallic tin. |

Nearly all stannous compounds are colorless; the oxide (as already
mentioned) is black and the sulfide dark brown.

252 HYDROGEN SULFIDE GROUP

REACTIONS IN THE WET WAY

1. Potassium and Sodium Hydroxides produce a white precipi-
tate of gelatinous stannous hydroxide,

Snt+++20H- > Sn(OH)o,

which is readily soluble in an excess of the precipitant, forming alkali
stannite: |

Sn(OH)2+OH~ — H20+H8n0-2_,

The hydroxide is also readily soluble in hydrochloric acid.

The alkaline solution of an alkali stannate is often changed to
brownish black or black (particularly on warming, or when very con-
centrated caustic alkali is used), owing to the separation of either
metallic tin or stannous oxide (ef. p. 170):

From dilute potassium hydroxide solutions there gradually sepa- .

rates on standing, or more rapidly on heating, the black monoxide,
HSnO2 — OH +Sn0O;

and from quite concentrated alkali the precipitate is almost wholly
black metallic tin:
2HSnO2 —SnO3 +H20+Sn.

2. Ammonia and Alkali Carbonates precipitate the white hydrox-.
ide, which is not absolutely insoluble in an excess of the precipitant: —

Sn**+20H- — Sn(OH)>;
Snt*+C0O37+H20 — COe 7 +$n(OH)o.

A large quantity of tartaric acid, more than is required in the
case of antimony, prevents the precipitation.

3. Hydrogen Sulfide produces (in solutions which are not too
acid) a brown precipitate of stannous sulfide,

Snt++H.S > 2H*+SnS,

readily soluble in strong hydrochloric acid; therefore no stannous
sulfide is precipitated if the solution is very acid. After diluting
a strongly acid solution with water, however, stannous sulfide is
completely precipitated on saturating the solution with hydrogen

sulfide gas.
Stannous sulfide is insoluble in ammonia and ammonium carbon-

ate (difference from arsenic); also in colorless ammonium sulfide (dif-

. * f
~~). os ae

TIN 253

ference from arsenic and antimony); but is readily soluble in yellow
ammonium sulfide, forming ammonium thiostannate:

SnS+ (NH4)2Se = (NH4)2SnSs3.

If the solution of ammonium sulfostannate is acidified with any
acid, yellow stannic sulfide is precipitated:

SnS3 +2H ex HS T +SnSo.

. 4. Mercuric Chloride produces in solutions of stannous salts a
white precipitate of mercurous chloride:

2HgCle+Sn* *t > Snt*+++2Cl-+H¢g.Clo.

But if the stannous chloride is present in excess, the mercurous
chloride will be reduced to gray mercury:

Hg2Cle+Sn** — Sn*t*+++2Cl-+2Hg¢g.

5. The Gold Test is much more sensitive. If to a solution of
gold chloride a solution containing a trace of stannous chloride is added,
finely divided metallic gold will be precipitated,

2Aut *++38ntt — 3Sntt+t+42Au,

which appears brown by transmitted light, and bluish-green by reflected
light.

6. Metallic Zinc precipitates tin from both stannous and stannic
solutions as a spongy mass, which adheres to the zinc:

Sn? ?-Zn = Zn t-+-Sn:

The finely-divided, spongy metal is easily soluble in strong hydrochloric
acid; the experiment must not, therefore, be made in strongly acid solutions,
as the tin is loosened from the zinc by the violent evolution of hydrogen and is
redissolved by the acid. The test is best made by adding a drop of the (not
too acid) solution to a piece of platinum foil, and then placing a piece of bright
zinc so that it comes in contact with both the golution and the platinum.
The tin is precipitated partly on the zinc and partly on the platinum,* in
the form of a gray stain, which disappears from the latter as soon as the zine
is removed, provided the solution is still acid (difference from antimony),
If the zinc is kept in contact with the acid until the evolution of hydrogen ceases,
the tin stain will not disappear from the platinum, because all the acid has been
used up. On adding a few drops of concentrated hydrochloric acid to the
platinum, the stain quickly disappears with an evolution of hydrogen. The
reason why the tin is deposited on the platinum notwithstanding the pres-

*In weakly acid solutions tin is precipitated chiefly on the zinc; in strongly
acid solutions, chiefly on the platinum.

254 - HYDROGEN SULFIDE GROUP

ence of acid is that a galvanic current is formed by the contact of the zine
with the platinum, which flows from the zine to the platinum; the platinum
thus serves as a cathode, and the tin is deposited upon it. On removing the
zinc the current stops and the stain disappears.

7. Metallic Lead, as is evident from its position in the electro-
motive series (cf. p. 41), will reduce tin to the metallic state, but the
reaction will stop as soon as the concentration of Pb** becomes nearly
as large as that of Sn**; when the concentration of Pb** is greater
than that of Sn** the reaction will take place in the reverse direction.
It requires a relatively low concentration of H* to stop both of these
reactions, as both tin and lead are higher in the series than hydrogen.
In acid solution, however, metallic lead will reduce Sn***+* to Snt*
(cf. p. 44).

Stannic Compounds

The stannic compounds (which are all colorless, with the excep-
tion of the yellow sulfide, SnSz2) cannot be obtained by the solution
of the oxide, SnOz, from which they are derived, because the oxide
is attacked with difficulty by acids. They are obtained indirectly
from metallic tin or from stannous compounds.

The simple stannic compounds are all, more or less readily, com-
pletely hydrolyzed by water, so that the analyst almost never meets
with them. The nitrate, Sn(NOs)4, and the sulfate, Sn(SO4)o, are
quickly decomposed in the cold, into acid and stannic hydroxide.
The halogen compounds are more stable, and are decomposed only
by boiling with considerable water. For the following reactions,
therefore, we will assume that we have a solution of stannic chloride
to work with. |

Stannic chloride is a colorless liquid, which fumes in the air and
boils at 120° C. On adding a little water it solidifies, forming crystals
of monoclinic hydrates, SnCl4a+3He2O, SnCl4a+5H20, SnCl4a+8H20,
of which the salt with 5H2O is used commercially as a mordant in
dyeing. ;

On adding more water to these hydrates they dissolve, forming
a clear solution, which on boiling (the freshly-prepared, dilute solution)
gradually becomes turbid, owing to the precipitation of voluminous
stannic hydroxide:

SnCl,+4HOH = 4HCl+Sn(OH)4q.
If the solution is very dilute it becomes turbid in the cold. The

stannic acid thus formed is not precipitated quantitatively, either
in the cold or on boiling, because a tonsiderable amount remains in

TIN 255

the hydrosole form. By ‘salting out” the hot solution (best with
ammonium nitrate), the stannic acid may be completely precipitated.
A solution of stannic chloride can be most readily obtained for
analytical purposes by chlorinating or brominating a solution of
stannous chloride.
On ‘adding chlorine to a solution of stannous chloride, stannic
chloride is formed in the cold:

SnCls 4. Cle = SnCl..

As, however, chlorine is colorless in a dilute solution, it is difficult
to tell when the oxidation is complete; it is more easily ascertained
if bromine is used.

On adding strong bromine water, drop by drop with constant
stirring, to a solution of stannous chloride, the brown color will dis-
appear as long as any stannous chloride remains unchanged, and the
solution becomes colored by the bromine only when the oxidation is
complete. ‘The solution then contains a mixture of stannic chloride
and stannic bromide:

2SnClo+2Bre=SnCl+SnBra.

Just as platinum tetrachloride combines with hydrochloric acid
to form chloroplatinic acid, so tin tetrachloride unites with hydro-
chloric acid, forming chlorostannic acid, HeSnCle, and yields, like
the former, beautifully crystalline, easily soluble salts with the alka-
lies, of which the ammonium salt (NH4)2[SnCl¢] is an article of com-
merce, being known as “ pink salt.’ The above-mentioned stannic
chloride, SnCl4, is sometimes designated as a-stannic chloride, to
distinguish it from a compound (which we shall soon study) known
as 6-stannic chloride (stannyl chloride).

Reactions of the a-Stannic Compounds

1. Hydrochloric and Sulfuric Acids produce in moderately con-
centrated solutions of stannic chloride no precipitation, even on long
standing (difference from #-stannic compounds). In very dilute
sulfuric acid solutions a precipitate of basic sulfate is sometimes
obtained. In very dilute hydrochloric acid solutions, also, a slight
turbidity is often formed, which increases on boiling the solution:

SnCly+4HOH = 4HCI+Sn(OH)s.

2. Potassium and Sodium Sulfates produce no precipitation in
the cold (difference from stannyl chloride), but on boiling all the tin
is precipitated as hydroxide.

256 ~ HYDROGEN SULFIDE GROUP

3. Potassium or Sodium Hydroxide.—On adding caustic alkali
to a solution of a stannic salt, a voluminous, gelatinous, white pre-
cipitate is obtained:

SnCli +40H- — 4CIl--+Sn(OH),.

The precipitate has the above formula when dried in the air, and
the formula He2SnOs if dried over sulfuric acid.

The precipitate dissolves readily in an excess of alkali hydroxide,
forming salts which are not derived from either of the above com-
pounds, but from He-[Sn(OH).6|, which has itself never been isolated:

(SnOH),+2KOH = K2[Sn(OH)<].

The hydroxide dissolves in ammonia also, but only in the absence
of ammonia salts.

By dissolving in alkali, stannic hydroxide behaves as an acid,
and according to Bellucci and Parravano,* the hexahydroxystannic
acid stands in the same relation to chlorostannic acid as hexahydroxy-
platinic acid to chloroplatinic acid:

Ho2[PtCle] Ho2[SnCle]
H.[Pt(OH)c] H.2[Sn(OH)g].

The salts of hexahydroxystannic acid are designated briefly as
stannates, or a-stannates, to distinguish them from the 6-stannates
or metastannates, which are derived from the polymer (H2SnQOs3)5
(see p. 257).

The ready solubility of e-stannic acid in cold dilute mineral acids
is very characteristic. It dissolves promptly in hydrochloric, nitric,
and sulfuric acids, behaving, in this respect, as a base. By boiling the
dilute acid solution (particularly the sulfuric acid solution) stannic
acid is reprecipitated, which is soluble in cold dilute acids provided
the boiling has not been continued too long. In the latter case the
6-stannic acid is formed, which is insoluble in dilute acids.

4. Potassium Carbonate precipitates stannic acid from stannic
chloride solutions; the precipitate is completely soluble in an excess
of the reagent (difference from 6-stannic acid):

SnCly -- 2K2CO3 a 2H2O =4KCl ao 2CO2 ad Sn(OH),.

5. Sodium Carbonate behaves similarly, but the precipitate is
not so easily soluble in an excess.
6. Ammonia precipitates stannic acid from a solution of stannic

* Z. anorg. Chem., 45 (1905), p. 156.

TIN | 257

chloride; tartaric acid prevents the precipitation (difference from
B-stannic acid).

Reactions of the 6-Stannic Compounds
(Metastannates)

By the oxidation of metallic tin with hot nitric acid of sp.gr. 1.3
stannic nitrate is first formed, which, by boiling with water, is com-
pletely hydrolyzed, forming nitric acid and metastannic acid.

Metastannic acid is a white powder insoluble in nitric acid, and
when dried over sulfuric acid has the formula H2SnO3. This is of the
same empirical composition as the hydroxide precipitated by treating
a stannic chloride solution with alkalies, though differing essentially
from it in many reactions.

While the a-stannic acid (as already mentioned) is easily soluble .
in dilute mineral acids, the 6-stannic acid is almost insoluble therein.

1. If the B-stannic acid is treated for a short time with concen-
trated hydrochloric acid, a chloride is formed which a insoluble in
hydrochloric acid, but readily soluble in water. The solution con-
tains the so-called 8-stannic chloride (though the designation stannyl
chloride would be more suitable) of the composition, Sn5O5Cle(OH)s.*

2. On treating the aqueous solution of stannyl chloride with
hydrochloric acid, almost all the tin is reprecipitated in the form of
a highly chlorinated compound of the composition Sn505ClufOH)6+
4H20.t ,

3. If an aqueous solution of stannyl chloride is heated to boiling,
almost all the tin is precipitated as 6-stannic acid, which is insoluble
in dilute acids.

This differing behavior of the two acids, as well as of the two
chlorides, can be explained as follows: Silicic acid, which is closely
related to stannic acid, exists in innumerable silicates.in different
polymeri: forms. Thus, with the minerals of the pyroxene and
amphibole groups, wollastonite, CaSiOs, is a derivative of ordinary
meta silicic acid and tremolite, CaMgs3Si4Oi2, is a derivative of
' (He2SiO3)4.f It is highly probable that the stannic acid can exist
in analogous polymers. One of these polymers apparently corresponds
to the composition (H2SnQs3)s.

If such a compound is treated with hydrochloric acid, the hydroxyl

*R. ENGEL, Chem. Zig., 1897, pp. 309 and 859.
+ Weser, Jahresber., 1869, 244, and Pogg. Ann., 122, 358.
t Grorn, Tabellarische Uebersicht d. Min., 1898, p. 148.

258 HYDROGEN SULFIDE GROUP

groups will, first, be replaced by chlorine, and a compound will be
obtained containing tin, oxygen, and chlorine, e.g., Sns5O5Clro.

This hypothetical compound of the 6-stannic acid is hydrolyzed,
forming different chlorides of varying solubilities. Thus R. Engel
found that the chloride SnsOs5Cle(OH)g is soluble in water; and
Weber showed that from an aqueous solution of the latter, hydro-
chloric acid precipitates the compound Sn5O5Cl4(OH)6+4H20.

The reaction which takes place on dissolving the §-stannie acid
in hydrochloric acid and water may be expressed satisfactorily by
the following equations:

Sn;05(OH)10+10HCl=10H20+Sn;05Clio (insoluble in HCl);

8-stannic acid

or
Metastannic acid

S$n505Clio+8H20 — 8HCI+S8n505Cle(OH)s (soluble in water) ;
Sn505Cle(OH)s+2HCl @ 2H20+S$n505Cl4(OH)¢ (insoluble in HCl).

On boiling the aqueous solution, complete hydrolysis takes place:
Sn505Clo(OH)s+2HOH =2HCI+Sn;505(OH) 0.

If the 6-stannic acid is treated for a long time with concentrated
hydrochloric acid, the Sn5Os5 group is finally broken down, and the
tin goes into solution in the form of ordinary a-stannic chloride:

$n50;5(OH)19+20HCI = 15H20+5SnCly.

Further reactions of stannyl chloride (6-stannic chloride) are:

4. Sulfuric Acid precipitates from solutions of stannyl chloride
white stannyl sulfate, which on being washed with water is completely
changed to $-stannic acid (difference from a-stannic chloride).

5. Potassium and Sodium Sulfates cause the same reaction as
sulfuric acid.

6. Potassium Hydroxide throws down in solutions of stannyl
chloride a voluminous precipitate of 6-stannic acid, which does not
dissolve in an excess of the concentrated precipitant, but forms a
6-stannate easily soluble in dilute caustic potash solution:

(a) Sns05Clo(OH)s+2KOH =2KCl+Sn;0;(OH) 10;

B-stannic acid

(b) Sn505(OH) 10 +2KOH = 2H20+8n505(OK)2(OH)s.

potassium #-stannate

By long treatment of the potassium #-stannate with concentrated
caustic potash, it gradually goes into solution, forming a-potassium

TIN 259

stannate. This change takes place more readily by fusing 6-stannic
acid with solid potassium hydroxide in a silver crucible.

If a dilute solution of a mineral acid is added to the 6-potassium
stannate, a voluminous precipitate is formed, consisting partly of
B-stannie acid (insoluble in an excess of mineral acids) and partly
of a-stannic acid (readily soluble in an excess of the acid). ‘The latter
compound is formed when a very concentrated solution of caustic
potash was used in forming the potassium salt.

7. Ammonia also precipitates 6-stannic acid, even in the pres-
ence of tartaric acid (difference from a-stannic chloride).

As we have seen, the a-compounds may be readily changed into
B-compounds and conversely. The dilute aqueous solutions of the
a-compounds are gradually changed, at the ordinary temperature, into
B-compounds, but more quickly on boiling; thus stannic chloride changes
to stannyl chloride:

58nCl4+13H20 =18HC1+Sns505Cle(OH)s.

The B-compounds are changed into a-compounds by boiling with
oncentrated hydrochloric acid or with concentrated caustic potash.

8. Hydrogen Sulfide precipitates (from not too acid solutions)
yellow stannic sulfide from both the.a- and the 6-compounds:

SnCl,+2Has a= 4HCl+- SnS2;
Sn505Cl2(OH)g+10H2S @ 2HCl1+13H20+5SnSo.

From £-stannic solutions hydrogen sulfide produces a precipitate, but
very slowly, the SnS, remaining largely in the hydrosole form. By the addition
of salts it is coagulated, and separates out in a flocculent form, usually mixed
with g-stannic acid. If the stannyl chloride solution is heated on the water-
bath in a pressure flask, the tin is quickly precipitated as greenish-yel'ow
sulfide.

Stannic sulfide is soluble in hydrochloric acid; hydrogen sulfide
will cause no precipitation, therefore, if the solution is very acid. If
such a solution, saturated with hydrogen sulfide, is largely diluted,
the sulfide will precipitate out.

The yellow sulfide is the thio-anhydride of the thio-acid; it dis-
solves, therefore, in alkali sulfide, forming salts soluble in water:

SnS2+ (NH4) 28 — (NH4)2Sn8sz.
Acids precipitate from such a solution the yellow stannic sulfide

SnS37-+2H* — HeS+SnSo.

260 HYDROGEN SULFIDE GROUP

The sulfide is insoluble in ammonia and ammonium carbonate
(difference from arsenic). By means of concentrated nitric acid it
is easily oxidized to 8-stannic acid; or by roasting in the air it can
be completely changed to tin dioxide.

The sulfide obtained in the dry way, known as “ mosaic gold,”
is not attacked by nitric acid, and is also insoluble in alkali sulfides.
It dissolves on treatment with aqua regia, forming stannic chloride
with separation of sulfur. It is most readily brought into solution
by fusing with sodium carbonate and sulfur (see below).

9. Mercuric Chloride produces no precipitation in solutions of —
stannic salts.

Tin dioxide as it occurs in nature, and the artificially produced oxide
after strong ignition, are both insoluble in all acids. They can be brought
into solution by the following methods: !

1. Fusion with sodium carbonate and sulfur;
2. Fusion with caustic potash or soda;

3. Fusion with potassium cyanide;

4. Reduction by hydrogen at a high heat.

(1) Fusion with Sodium Carbonate and Sulfur.—Place the dry substance
in a small porcelain crucible, mix with six times as much calcined sodium
carbonate and sulfur (equal parts mixed together), cover the crucible, and heat
over a small flame until the excess of sulfur has distilled off and burned. This
operation requires about twenty minutes. Allow the crucible to cool and
then treat its contents with hot water, and filter if necessary: .

2S8n0, +2Na.CO; +95 =3S0, T +2Na.Sn8;+2CO, tT °

If iron, lead, copper, or any other metal that forms a sulfide insoluble in
water and ammonium polysulfide is present, it remains undissolved as sulfide,
and is separated from the tin by filtration.

(2) Fusion with Sodium Hydroxide.—Melt the sodium hydroxide in a silver
crucible, placed within a porcelain crucible to protect the silver from the
injurious action of the gas flame. When the fusion has become quiet, owing
to the expulsion of all the excess water, cool somewhat and add the finely-
powdered, dry tin dioxide. Heat until the fusion is clear. After cooling
dissolve the melt in water:

SnO, +- 2Na0H = NaSnQ; + H.0.

Stannic oxide is not completely attacked by fusion with sodium or potassium
carbonate.

(3) Fusion with Potassium Cyanide-—Melt some potassium cyanide in a
porcelain crucible, add the powdered stannic oxide and fuse the mixture until
‘the separated tin has melted together:

Sn0.+2KCN =2KCNO-+Sn.

After cooling, extract the mass with water, filter off the tin, flatten it into
foil and dissolve it in concentrated hydrochloric acid.
(4) Reduction in a Slream of Hydrogen.—Place the substance in a platinum

‘9

a ae

GOLD 261

boat and insert the latter in a glass tube, open at both ends, which is made
of difficultly fusible glass. Pass hydrogen through the tube in the cold until
the air has been driven out, and then heat to dull redness until no more water

is formed:
SnO,.+4H =2H.0+S8n.

Finally cool and dissolve the metal in hydrochloric acid.

REACTIONS IN THE DRY WAY

If a tin salt is heated with soda (or potassium cyanide) on char-
coal, usually only a small malleable button is obtained, which, on
taking away the flame, is immediately covered with a white coating
of oxide. This can be observed when the flame is allowed to play
upon the fusion. If the product is crushed in an agate mortar, a
small flake of metallic tin is obtained, which can be distinguished
from silver and lead by its insolubility in nitric acid, and by its solu-
bility in hydrochloric acid. This reaction is particularly suitable
for the charcoal-stick test. The borax bead which has been colored
pale blue by copper becomes a transparent ruby red in the reducing
flame if a trace of tin is added. This is a very sensitive reaction.

GOLD, Au. At. Wt. 197.2
Sp. Gr.=19.33. M. Pt. =1063° C.

Occurrence.—Gold usually occurs native in quartz and in river
sands; also as telluride of gold in sylvanite, (AuAg)2Tes, and in
- nagyagite, (PbAu)2(TeSSb)3, and is found in small amounts in many
pyrite and other sulfide ores.

Metallic gold is of a yellow color and melts without being oxidized.
It is the most ductile of all metals, and may be hammered into
exceedingly thin leaves, which are transparent, with a bluish-green
color.

Commercial gold is usually alloyed with copper, silver, or with
both metals.. Pure gold is designated as 24-carat gold or !9°/j000
fine. Fourteen-carat gold contains 14 parts of gold to 10 parts of
other metal and 18-carat gold contains 18 parts of gold and 6 parts
of alloy; the former, therefore, contains 58.3 per cent and the latter
75 per cent of pure gold.

The proper solvent for gold is aqua regia, but it is also soluble in
bromine and chlorine water, forming a trihalogen compound:

2Au+2HNO3+6HCl=4H20+2NO T +2AuCls,
Au+3Br=AuBrs.

262 HYDROGEN SULFIDE GROUP

Gold is not» attacked by minerals. It forms two oxides: aurous
oxide, Au2O; and auric oxide, AugQ3.

Both of these are exceedingly unstable; on gentle ignition they
lose oxygen and are changed to metal (a property common to all
“noble ”’ metals).

All gold salts are unstable; even the most stable salt of all, the
chloride, AuCls, is changed by gentle ignition into yellowish-white
aurous chloride, AuCl:

AuCls = AuCl +Cle.

On stronger ignition the last atom of chlorine is lost, and the yellow
metal itself is left behind.

Aurous chloride, AuCl, is insoluble in water, but on being boiled
with water for some time, or very gradually in the cold, it is changed
to auric chloride, with deposition of metal:

3AuCl= AuCl3+2Au.

The solution obtained by dissolving gold in aqua regia always
contains auric chloride, so that only the reactions of aurie com-
pounds are of interest to the analytical chemist. Auric chloride
unites with hydrochloric acid, formiug chlorauric acid,

AuCl3-+HCl=H[AuCly],

which yields beautifully crystalline salts.

Gold chloride is soluble in ether and can be extracted from its
aqueous solutions by means of this solvent. :

Auric salts are mostly yellow and readily soluble in water. The
sulfide is black and soluble only in aqua regia.

REACTIONS IN THE WET WAY

A solution of chloraurie acid, H{[AuCl4], should be used for these
reactions.

1. Potassium or Sodium Hydroxide.—If caustic alkali is cau-
tiously added to a concentrated gold solution, a reddish-brown, volu-
minous precipitate of auric hydroxide is obtained, which looks exactly
like ferric hydroxide. If more caustic alkali is added, however, the
gold hydroxide redissolves, forming alkali aurate:

Autt+*++30H-— Au(OH)3; Au(OH)3+OH~ — 2H20+[AuO,]-.

If the bright-yellow solution of potassium aurate is carefully acidi-
fied with nitric acid, a precipitate of reddish-brown auric acid is thrown

GOLD 263

down, which is soluble in nitric acid, but is reprecipitated, for the
most part, by diluting and boiling.

As a rule, potassium hydroxide yields no precipitate in solutions
of gold chloride, because the gold solution is usually so dilute that
the amount of alkali added is sufficient to form the aurate at once.

2. Ammonia throws down a yellowish mixture of gold imino-
chloride, Au(NH)Cl, and gold iminoamide, Au(NH)NHa2,

AuCl3+3NH3 — 2NH4Cl+Au(NH)Cl1,
AuCl3-+5NH3 — 3NH4Cl+Au(NH)NHa2,

which explodes in a dry condition on warming or by concussion (ful-
minating gold.*)

The most important reactions for the detection of gold are those which
depend upon the extreme readiness with which the auric compounds are
reduced. Auric compounds are strong oxidizing agents.

3. Ferrous Salts precipitate at ordinary temperatures from neutral
or acid solutions all the gold as a brown powder (difference from
platinum): :
Aut+++43Fet+— 3Fet+++Au.

4. Oxalic Acid precipitates all of the gold in the cold, but more
quickly on warming (difference from platinum):

2Aut 7 +41 3H C204 nex. 6Ht+6CO2 T +2Au.

The presence of considerable acid prevents this reaction.
5. Arsine and Stibine precipitate gold completely:

2Aut*++AsH3+3H20 — H3AsO3+6H*+2Au.
6. Sulfurous Acid reduces gold solutions:
2Aut**+3S802+6H20 — 12H++3S804- +2Au.
7. Stannous Chloride causes the following reaction to take place:T

2Autt+++3Snt+— 3Snt*t*+2Au.

If the solution tested is strongly acid with hydrochloric acid, the precipi-
tate is pure gold and has the characteristic dark-brown color of the finely
divided metal. In very dilute weakly acid solutions the so-called purple
of Cassius is thrown down, which consists of colloidal gold and tin hydroxide.f

Purple of Cassius is soluble in ammonia and in dilute caustic potash solu-

* Rascuia, Ann. Chem. Pharm. 235 (1886), 325.
+ THEopor Doérina, Chem. Centralbl., 1900, I, p. 735.
t Zstamonpy, Ann. Chem. Pharm., 301 (1898) 361.

264 HYDROGEN SULFIDE GROUP

tion, forming reddish liquids. These solutions when cold remain clear for a
long time and can even be boiled wlthout decomposition. As the solution
is concentrated a flocculent precipitate is formed which will dissolve on the
addition of more ammonia.

The brown coloration can be distinctly seen if 0.3 mg. gold is dissolved

in 100 cc. of the solution; if less than 0.3 mg. of gold is present, only a yellow-—

ish coloration is obtained.

8. Hydrogen Peroxide* in alkaline solution immediately pre-
cipitates the gold as finely divided metal:

2Au***+3H202+60H — 6H20+302 fT +2Au.

The precipitated metal appears brownish-black by reflected light,
but bluish-green by transmitted light; 0.03 mg. gold in 10 ce. of liquid
suffices to give a reddish coloration with a bluish shimmer.

9. Zinc.—The following gold test is very sensitive: To a few
drops of a dilute gold solution add a few drops of arsenic acid, two
or three drops of ferric chloride solution, and two to three drops of
hydrochloric acid. Dilute the mixture to 100 cc., and drop in a piece
of zinc. Around the zinc the solution assumes a purple color, which,
by moving the zinc in the solution, is disseminated through it, making
it appear pink or purple.

If the solution contains 0.03 mg. of gold, within fifteen minutes a
beautiful reddish color will be noticed.

Besides the above reagents many others, such as formaldehyde
in the presence of alkali, hydrazine sulfate, etc., are capable of reducing
gold from its solutions.

10. Hydrogen Sulfide iia Bc in the cold, black gold disul-
fide from gold solutions:

2Aut ¥% +43HS —J 6Ht+ S +AuoSo.

Gold disulfide is insoluble in acids; but is readily soluble in aqua
“regia, forming auric chloride, AuCls.

The disulfide is difficultly soluble in ammonium sulfide, but more |

readily soluble in potassium polysulfide, forming a sulfo salt:
Aur2S2+ K2Se = 2K AuSo.

From this solution hydrochloric acid precipitates a yellowish-
brown sulfide:
2KAuSe+2H* ae 2K*+ H2S +Au2S3(?).

* Vanino and SEEMANN, Ber., 1899, 1968.
+ Pharm. Chem. Centralbl., 27, 321.

Soe

— of a. eT

GOLD 265

From a hot solution, hydrogen sulfide precipitates brown, metallic
gold:
S8Au’ +++3H2S+12H20 — 30H*+3804"- + 8Au.

The finely divided metallic gold is soluble in hot potassium or
sodium polysulfide, forming a thio salt:

2Au+ KoS4 <—F 2KAuSo.

On account of its softness, gold is always alloyed with silver and copper
when used for coins or for jewelry. If such an alloy is treated with nitric
acid, the copper and silver are dissc!ved and the gold usually remains as a
brownish powder. ‘To prove the presence of gold, filter through a small filter,
dry the filter, roll up the paper and wrap a platinum wire around it. Set
fire to the paper and allow it to burn quietly.. The ash must not be too
strongly heated, for the gold would then melt and alloy with the platinum
wire. Melt the ash with a little sodium carbonate on a charcoal stick; a gold
button forms with the characteristic yellow color. The gold button can be
pressed into a leaf in the agate mortar, transferred to a watch-glass, and
dissolved in a little aqua regia, forming auric chloride. Carefully evaporate
the solution to dryness, dissolve the residue in a little water, add a dilute solu-
tion of stannous chloride, and the presence of gold is shown by the formation
of the purple of Cassius. The hydrogen peroxide or zinc tests are still more
delicate (see p. 264).

If it is a question of detecting very small amounts of gold (as in the case
of many copper coins), the above method is unsuitable. In such cases,
extract the gold and silver by melting with lead, and remove the lead by cupel-
lation. Melt 5 to 10 gms. of the auriferous copper (or more in some cases)
with 120 gms. of pure lead in a flat dish of infusible stone (a scorifying-dish),
in a muffle with access of air. The copper and a part of the lead are oxidized,
and the oxide unites’ with the silica of the dish to form a readily fusible slag,
which eventually covers the unaffected lead and the dissolved silver and
gold. This operation is known as scorification. When this point is reached,
pour the molten mass into an iron scorification pan, previously well chalked.
As soon as the mass becomes cool, remove it from the pan, and hammer the
slag from the enclosed lead button and weigh the latter. Then put it on a
cupel (a sort of crucible made of bone ash), of about the same weight as the
lead button or a little heavier. Place the cupel in the muffle and again heat
with ready access of air. The lead melts and is oxidized; the resulting lead
oxide melts at 980° C. and is partly absorbed by the porous eupel and partly
volatilized, leaving a kernel of silver and gold behind. Tlatten the metallic
kernel to foil and treat it with 6-normal nitric acid, which dissolves the silver,
leaving the gold, usually in the form of powder. Filter off the gold, dry, and
melt it, as above described, upon the charcoal stick. If the alloy of gold and
silver (obtained after cupellation) contains three parts of silver to one part of
gold, the gold remains after separation with nitric acid as a thin-as-paper, coher-
ent, brownish mass (which becomes hard on igniting), with the characteristic
gold color. If the proportion of silver to gold is greater than 3 : 1, the sepa-
ration by means of nitric acid will be complete and the gold will be left as a
powder, If the ratio of silver to gold is less than 3 : 1, the separation by means

266 HYDROGEN SULFIDE GROUP

of nitric acid is incomplete, and the gold residue usually appears yellow, and
still contains some silver. In‘this case, add more silver with 1 gm. of lead
and subject the mixture once more to cupellation, when the subsequent sepa-
ration by means of nitric acid will be complete.

In order to detect very small amounts of gold in ores, a similar procedure
is used.

If one does not possess a muffle furnace, the tiresome wet process must
be used. For example, if it is desired to detect the presence of gold in pyrites,
roast a large amount of the ore in the air until all the sulfur has been burned
off, then treat with bromine water, and allow to stand twelve hours. Filter
the solution (which now contains all of the gold as auric bromide) and boil
off the excess of bromine. Add ferrous sulfate and a little sulfuric acid, boil
the solution again and filter through a small filter. Wash the residue on the
filter, dry and then melt it on the charcoal stick. |

According to these methods a fractional part of a milligram of gold can be
detected with certainty.

REACTIONS IN THE DRY WAY

All compounds of gold, when heated with soda on the charcoal
stick, yield a malleable, metallic button, soluble only in aqua regia.
The solution in the latter reagent should be evaporated, the residue
dissolved in water and tested with stannous chloride, hydrogen per-
oxide, or zinc.

PLATINUM, Pt. At. Wt. 195.2
Sp. Gr. =21.48. M. Pt.=1755° C.

Occurrence.—Platinum is found free in nature, usually accom-
panied by the other so-called platinum metals.

Metallic platinum is grayish white; in a finely-divided state it is
grayish black. The metal is not ordinarily attacked by mineral acids; *
it dissolves in aqua regia, forming chloroplatinic acid, He[PtCle] (not
platinum chloride, PtCl4). If, however, the platinum is alloyed
' with silver, provided sufficient silver is present, it dissolves in nitric
acid, forming a yellow solution. Like tin, platinum forms two
oxides: platinum monoxide, PtO; and platinum dioxide, PtOc.

Both oxides may be obtained by the careful ignition of the cor-
responding hydroxides. They are exceedingly unstable, being decom-
posed by gentle ignition into metal and oxygen; all the remaining
platinum compounds behave similarly. |

The most important of the platinum compounds are the chlo-

* By boiling concentrated sulfuric acid in platinum dishes 10 cc. of acid will
dissolve 3.8 mg. of platinum in ten minutes. Le Roy and McCay, 8th Internatl.
Cong. App. Chem. 1, 351 (1912).

PLATINUM 267

rides. By dissolving platinum in aqua regia, chloroplatinic acid is
always obtained, from which the di- and tetrachlorides may be
derived; these unite with hydrochloric acid to form the complex acids

PtCl4a+2HCl = Ho2[PtCle] (Chloroplatinic acid—orange red crystals),

PtCle+2HCl=Hoe[PtCl4] (Chloroplatinous acid—known only in solu-
tion). |

The aqueous solution of chloroplatinic acid is yellowish orange;
a solution of chloroplatinous acid, containing the same quantity of
platinum, is dark brown.

The potassium and ammonium salts of chloroplatinous acid are
soluble in water; the corresponding salts of chloroplatinic acid are
difficultly soluble in water and insoluble in 75 per cent alcohol.

REACTIONS IN THE WET WAY

A solution of chloroplatinic acid should be used for these reactions.
1. Ammonium and Potassium Chlorides produce in concentrated
solutions yellow precipitates (cf. pp. 78 and 89):

[PtClg)"+2K+t > Ko[PtCle]; [PtCle]- +2NH4t > (NH4)2[PtClo];

Both salts are difficultly soluble in water, but practically insoluble
in 75 per cent alcohol and in concentrated solutions of potassium
and ammonium chlorides. This last property is utilized in separating
platinum from gold and other metals.

2. Alkali Iodides give a brownish-red coloration due to the forma-
tion of the less ionized [PtI¢]~-

[PtCle]-+6I- — [PtIe¢]—-+6ClI-.

3. Hydrogen Sulfide precipitates in the cold very slowly, but
quickly on warming, dark-brown platinum disulfide:

H.2[PtCle]+2H2S =6HCI1+PtSo.

Platinum sulfide is insoluble in mineral acids, but readily soluble
in aqua regia. It is difficultly soluble in alkali sulfides, but more
readily soluble in alkali polysulfides, forming a thio salt, which is
decomposable by acids, with precipitation of platinum sulfide.

4. Ferrous Salts do not reduce chloroplatinic acid in the pres-
ence of acids (difference from gold), but cause precipitation of all
the platinum (on warming) in a solution which has been neutralized
with sodium carbonate:

268 HYDROGEN SULFIDE GROUP

[PtCle]- +6CO37+4Fet ++6H20 — 6CO2 T +6CI-+4Fe(OH)3+Pt.

5. Oxalic Acid does not precipitate platinum (difference from
gold).

6. Formic Acid precipitates from neutral, boiling solutions all the .

platinum in the form of a black powder:
H2[PtCle] +2HCOeH =6HCI+2CO02+Pt.

Formic acid

An acid solution must be neutralized with sodium carbonate before
making this test.

7. Stannous Chloride reduces chloroplatinic acid to chloroplatin-
ous acid only, not to metal:

[PtCle]-+Snt + =Snt*+t+++2Cl-+[PtCh]-.

8. Glycerol and Sodium Hydroxide reduce chloroplatinic acid on
warming, with the separation of black, pulverulent metal:

C3Hs(OH)3+8[PtCle]> +16OH- — 18Cl +CO37 +C2047 +12H20 +8Pt.

9. Carbon Monoxide, on being passed into a solution of chloro-
platinic acid containing sodium acetate, colors the solution a beau-
tiful red owing to the formation of colloidal platinum (difference from

palladium). After standing some time all the platinum is precipitated a ,

as a black powder and the supernatant liquid is colorless:

H2[PtCle] +2CO+6C2H302°+2H20 -6Cl-+6HC2H302+ 2COe2T +Pt.

10. Zinc, cadmium, magnesium or aluminium reduces chloroplatinic

acid to metal:
H2[PtCls]+3Zn — 3ZnCle+He ft +Pt.

The precipitated metal is in such a finely divided condition that it
tends to run through the filter, especially on being washed with pure
water; by washing with salt solution a clear filtrate can be obtained.

11. Hydrazine Hydrochloride, NoH4-2HCl, readily reduces chloro- ; ’

platinic acid in ammoniacal solutions; the platinum is deposited
partly as mirror upon the sides of the vessel containing the solution.

(NHa4)2[PtCle]+N2H4-2HCI+6NH3 — 8NH4Cl+Ne J +Pt.

12. Formaldehyde in alkaline solutions precipitates the platinu
as extremely- finely-divided platinwm black:

He[PtCl6]| +HCHO+60H- — COz fT +6CI”-+5H20+Pt.

The precipitate may be washed with alkali salt solutions, but with pure ©

water a black colloidal solution of platinum is formed.

ee ee a ee ee ne en ee eee

PLATINUM 269

Preparation of Chloroplatinic Acid for Use as a Reagent

Since chloroplatinic acid is used not only for the qualitative detection of

- ammonium and potassium, but also for their quantitative separation, methods

for preparing a solution of this reagent will be described.

1. Preparation of Chloroplatinic Acid from Metallic Platinum.—Almost
all commercial platinum contains iridium; and although pure iridium is prac-
tically insoluble in aqua regia, it dissolves considerably in this reagent if it

is alloyed with platinum. Moreover, platinum forms with aqua regia not

only cholorplatinic acid, but also chloroplatinous acid (the most harmful
of all impurities for this reagent) and nitroso-platinic chloride, [PtCl](NO)s.
These facts must be borne in mind in preparing the reagent.

First of all, clean the strips of platinum by boiling them with concentrated

hydrochloric acid and washing with water. Then place the platinum in a

capacious flask, pour over it concentrated hydrochloric acid, and add nitric
acid little by little, with continuous, gentle heating on the water-bath. All
the platinum and some iridium are thus brought into solution, while some
of the latter metal usually remains undissolved as a black powder.

Decant the solution (without stopping to filter) into a porcelain evapo-
rating-dish and evaporate to sirupy consistency. Dissolve the residue in
water, add some sodium formate and sodium carbonate until the solution is
slightly alkaline. Heat the solution to boiling, which causes the precipitation
of the platinum and iridium in a few minutes as a black powder. This opera-
tion should be performed in a large evaporating-dish, on account of the violent
effervescence due to the escape of carbon dioxide. Pour off the supernatant
liquid and wash the residue several times with hydrochloric acid to remove
the sodium salt, and finally with water to remove the acid. Dry the powder,
which contains iridium and platinum (in the presence of one another, but
not alloyed together) and ignite it strongly in a porcelain crucible over the
blast-lamp (whereby the iridium is made insoluble in aqua regia), and
then weigh. Dissolve the ignited gray metal (at as low a temperature as
possible) in hydrochloric acid, with gradual addition of nitric acid. Consider-
able quantities of nitroso-platinic chloride are formed by this operation. On

evaporating with water, this compound is decomposed into chloroplatinie

acid, with evolution of oxides of nitrogen:
[PtCl.](NO).+H,0=NO, T +NO f +H,[PtCl,J.

As, however, a part of the NO. (or N.O,) remains in solution, some more
nitric and nitrous acids are formed by the action of water,

N20, + H,O = HNO; -}- HNO,,

which yields nitrosyl chloride with the hydrochloric acid present, and causes
the formation of more nitroso-platinic chloride.

It is necessary, therefore, to evaporate the solution alternately with hydro-
chloric acid and with water until no more nitrous fumes are given off. The
solution thus obtained always contains some chloroplatinous acid, and is
intensely brown. In order to change this last compound into chloroplatinic
acid, saturate the warm solution with chlorine gas (whereby its color becomes
much lighter) and then evaporate (at as low temperature as possible) till it

270 HYDROGEN SULFIDE GROUP

becomes of sirupy consistency. After cooling, the sirup crystallizes to a
yellowish-brown mass, which may be dissolved in a little cold water, and the
insoluble iridium filtered off.

If there is a considerable amount of the latter metal, ignite it in a pongelaan
crucible and weigh. The weight of the iridium should be deducted from the
prev ious weight of the mixture, in order to find out how much platinum re-
mains in solution.

Dilute the filtered solution with water until 100 ec. of the solution contains

10 gms. of platinum.
2. Preparation of Chloroplatinic Acid from Platinum Residues.—These
residues consist of potassium chloroplatinate and the alcoholic wash-waters.
By evaporating an alcoholic solution of chloroplatinic acid, chloroplatinous
acid and ethylene are formed, which yield ethylene platinous chloride: this last
compound gives no precipitation with potassium or ammonium salts:

H.[PtCl]+2C,.H;OH =CH;CHO+4HCI+C,H,[PtCl.]4+ H.0.
Alcohol Aldehyde

On evaporating an alcoholic solution of this soluble organic platinum
compound, it is changed into an insoluble powder which is explosive when dry,
insoluble in acids, and completely decomposed by strong ignition only.

To separate platinum, therefore, from platinum residues, first evaporate
the alcoholic solution to dryness, take up the residue in water, and pour the

solution into caustic soda (sp.gr. 1.2), to which 8 per cent of glycerol * has

been added. Heat the liquid to boiling, which causes the platinum to be pre-
cipitated as a heavy black powder,

FOO. GHP OCU HO = 36HCI+ 2CO.+2H.2C.0.+6Pt.

Glycerol Oxalic acid

Wash the powder first with water, then with hydrochloric acid, and finally
with water again. Dry, ignite (to destroy any of the organic compound),
weigh and transform, as before, into chloroplatinic aicd.

REACTIONS IN THE DRY WAY

All platinum compounds, when heated with soda on charcoal,
are reduced to the gray, spongy metal, which assumes a metallic luster
on being rubbed with a pestle in an agate mortar. It can be dis-
tinguished from gold by its color, and from lead, tin, and silver by its
infusilibity and insolubility.

Separation of Gold from Platinum

Precipitate the platinum with a solution of ammonium chloride;
filter and treat the solution with ferrous sulfate to precipitate the
gold.

*Z. anal. Chem., 28, 509.

a ee ca el i i

ANALYSIS OF GROUP II 271

Analysis of Group II

Gold and platinum are two well-known metals which really belong to this
group, but they are seldom present in large quantities in ordinary mineral
analysis and the detection of small quantities is so important that it is custom-
ary to test for them separately as has been indicated (cf. pp. 265 and 270).

The metals of the second group are precipitated as sulfides by means of
hydrogen sulfide in the presence of 0.3-normal mineral acid. The theory
governing the precipitation of sulfides was discussed on pp. 13, 47, 145, 186,
etc. In 0.3-normal mineral acid, the concentration of sulfur ions, formed by
the ionization of hydrogen sulfide, is made so small that the solubility products
of the sulfides of zine, nickel, cobalt, iron and manganese is not reached unless
large quantities of these elements are present in solution. In quantities up
to 0.5 gm. of metal, zine sulfide is the only one which is at all likely to pre-
cipitate in acid of this concentration, and then only when it is carried down
with a considerable quantity of some second-group metal. The solubility-
product of lead sulfide is about 4.2<107-28 and that of cadmium sulfide is
3.6X10~*9, while that of zine sulfide is 1.210-?%. Compared with mercuric
sulfide, with its solubility-product of 4.010753, lead sulfide is quite soluble,
but compared with zine sulfide, it is much less soluble. In precipitating these
sulfides there is some tendency for the sulfide precipitate to carry down
with it some of the metals of the succeeding groups partly by adsorption, but
this tendency is so slight that even with zinc the test is usually obtained in
the proper place when as much as 2 mgms. is originally present.

The solubility of the sulfides of the second group varies so greatly that it
is possible to separate them from one another by regulating the acid con-
centration so that enough sulfur ions are furnished by hydrogen sulfide to
precipitate one metal and not another. For the same reason, when hydrogen
sulfide is passed into the acid solution, the most insoluble sulfide is precipitated
first and very little, if any, of a more soluble sulfide is formed until the precip-
itation of the former is complete. In this respect the behavior of arsenic forms
an apparent exception; but this is due, as already pointed out (p. 230), to the
absence of an appreciable quantity of arsenic cations in the solution of an
arsenate. The order in which the metals are precipitated as sulfides from
cold solutions, as the acid concentration is progressively decreased, is approx-
imately as follows: arsenic, mercury and copper; antimony, bismuth and stan-
nic tin; cadmium, lead and stannous tin; zinc, iron, nickel, cobalt and man-
ganese. It will be noticed that zinc is placed a little in front of nickel and
cobalt in this arrangement. This is the proper order for the precipitation, but
after the sulfides have been formed a short time, nickel and cobalt sulfides
become less soluble than freshly-precipitated cadmium, lead or stannous
sulfide.

The theory of the solution of the sulfides has also been indicated (pp. 145,
181). The more soluble sulfides can be dissolved by merely increasing the con-
centration of the hydrogen ion. Thus the sulfides of antimony, tin, lead, cad-
mium, etc., can be dissolved by treatment with concentrated hydrochloric
acid.. Heating the solution to expel hydrogen sulfide hastens the dissolving,
but the most effective expedient is to add an oxidizing agent, which oxidizes the
sulfur ions as fast as they are formed in solution. Thus even mercuric and
arsenic sulfides will dissolve readily on being treated with aqua regia or with
bromine water,

272 HYDROGEN SULFIDE GROUP

The second group of metals may be divided into two smaller sub-groups.
The first of these contains mercury, lead, bismuth, copper and cadmium and
represents those elements whose sulfides are not dissolved readily by ammonium
polysulfide. Most of these sulfides show a slight tendency to dissolve in the
polysulfide and the tendency is most marked with the sulfides of mercury

and copper. From 5 to 10 mgms. of CuS and 0.5 to 1.0 mgms. of. HgS may dis-
solve in the polysulfide when the original substance contains considerable of —
these elements. None of these sulfides dissolve to any extent in ammonium ~

monosulfide.

The other subgroup contains arsenic, antimony and tin. The sulfides
of these elements dissolve readily in 6-normal ammonium polysulfide solution.
It is easy to dissolve 0.5 gm. of arsenic either as As»S; or as As»S;, 0.5 gm. of
antimony as Sb.S; or 0.5 gm. of tin as SnS, in ether ammonium monosulfide
or the polysulfide. Scarcely any SnS and only about 0.1 gm. of antimony as
Sb.S; will dissolve in 25 ce. of 6-normal ammonium monosulfide, but as much
as.0.5 gm. of either SnS or Sb.S; will dissolve in the same quantity of am-
monium polysulfide. Ammonium monosulfide dissolves scarcely any of the
sulfides of the other subgroup, so that when the antimony and tin are known
to be present in the higher state of oxidation, it is advisable to use ammonium
sulfide rather than the polysulfide for the separation of the two subgroups.

TasBLeE VII.—SEPARATION OF THE COPPER AND TIN GROUPS

Solution may contain: Hgt+, Pbtt, Bitt++, Cutt, Cdtt, Astt+t, AsO;=, AsOg,
Sbt+++, Sbht++++, SbO=, Sntt, Snt++** and Groups III, IV and V. Make the
solution 0.3-normal with HCl and saturate with H»S. Filter and examine the filtrate
for Groups IIIT, IV and V. Treat the precipitate with (NH4).Sz. (1)

Residue: HgS, PbS, Bi.S;, | Solution: [AsS,]=, [SbS,]=, [Sn83]>.
CuS, CdS. Add HCl. (3)
Examine by Table VIII. (2)

Precipitate: As.S;, SbS;, | Filtrate: NH,Cl.
Sn82: Reject.
Examine by Table IX. (3)

PROCEDURE

1. Add to the neutral solution 5 ec. of 6-normal hydrochloric acid and
saturate with hydrogen sulfide in the cold. Filter promptly and wash with
hot water containing hydrogen sulfide. Heat the filtrate to boiling and pass
hydrogen sulfide into it for ten minutes to make sure that arsenic will

be precipitated. If a yellow precipitate is formed slowly by this treatment, ~

filter it off through a new filter and evaporate the filtrate nearly to dryness.
Then add 10 ce. of 12-normal hydrochloric acid, saturate the cold solution
with hydrogen sulfide, heat nearly to boiling and again introduce hydrogen
sulfide. Cool, dilute to 100 cc. and saturate again with hydrogen sulfide.
Filter off the precipitate and examine the filtrate for the metals of Groups
III, IV and V (cf. p. 189 or 192).. This repeated treatment with hydrogen
sulfide is necessary when considerable arsenic is present (cf. p, 230).

Pe eee

ANALYSIS OF GROUP II 273

2. Transfer the precipitated sulfides to a porcelain dish and add about
10 cc. of ammonium polysulfide (or, better, ammonium monosulfide if it is
certain that the precipitate contains no stannous sulfide nor a large quantity
of antimony trisulfide). Cover the dish and warm gently for a short time with
frequent stirring. Dilute with a little water, filter and wash with hot water

containing a little ammonium sulfide. If the original residue was large and

the treatment with ammonium sulfide has evidently reduced its bulk, it is advis-

+ able to repeat the treatment with ammonium sulfide in order to make sure that
all of the arsenic, antimony and tin sulfides have been converted into soluble

thio salts. Examine the residue by the method outlined in Table VIII.

3. Dilute the ammonium sulfide extract with an equal volume of water,
add hydrochloric acid until the stirred solution is acid to litmus, heat to
boiling and filter as soon as the sulfides of arsenic, antimony and tin have
settled. Reject the filtrate and examine the precipitate by the method out-
lined in Table IX.

TaBLE VIII.—ANALysIs OF THE CopPpER GROUP.

Residue from Table VII: HgS, PbS, BiSs, CuS, CdS. Boil with 2-normal HNO3. (1)

Residue: HgS. | Solution: Pb++, Bit++, Cutt, Cdt++. Add 6-normal H2SOs,
Dissolve in evaporate, dilute and filter. (8).
aqua regia or
in bromine | precipitate: | Filtrate: Bit++, Cut+, Cd++.
water. Di-| ppso,. Add 6-normal NH,OH in excess. — (5)
oe a bil Dissolve in
ye | NAGC2H02 Precipitate: Solution:

PROCEDURE

1. Treat the residue from the ammonium sulfide treatment (Table VII)
with 2-normal nitric acid, boiling for a short time. Filter and wash with hot
dilute nitric acid. Examine the filtrate by (38).

2. Transfer the residue, and a part of the filter if necessary, to a porcelain
dish and digest with hot aqua regia, or warm with saturated bromine water.
Boil off the excess of chlorine or bromine, but do not evaporate to dryness on
account of the danger of losing some mercuric salt by volatilization. Dilute
somewhat and filter. Sometimes, when the elements of the copper group
are present in large quantity, the residue of sulfur (and filter paper) will con-

274 HYDROGEN SULFIDE GROUP

tain metastannic acid. To recover the tin, digest the residue with bromine
again if it is at all dark colored, filter and reject the filtrate. Then warm
the_residue with a little ammonium sulfide solution and add it to the solu-
tion obtained upon treatment of the original hydrogen sulfide precipitate with
ammonium polysulfide (Table VII). Test the solution obtained by the above
treatment with aqua regia, or bromine, for mercury by adding stannous chloride
solution. <A white precipitate of H. gnCle or Hg.Br2 shows that mercury is present.
The precipitate turns gray when an excess of stannous chloride is added
(cf. p. 200).

3. To the filtrate obtained in (1) add 10 cc. of 6-normal sulfuric acid and
evaporate in a porcelain dish until white fumes of sulfuric acid are evolved.
The purpose of this treatment is to expel all nitric acid, in which lead sulfate is
somewhat soluble. Cooland pour into 10 ce. of water, rinsing out the dish with
a little cold water. Stir well and let the solution stand five minutes but not much
longer. Filter and treat the fitrate by (5). When much bismuth is present
some basic bismuth sulfate is often precipitated with the lead sulfate, but the
greater part of the bismuth will remain in solution, particularly if the water
is kept cold. The basic bismuth sulfate, (BiO)»SO,, is more coarsely crystalline
than lead sulfate. When such a precipitate is obtained, treat it with 10 ce.
of hot 6-normal hydrochloric acid, add 10 cc. of 6-normal sulfuric acid and evapo-
rate as-above. The precipitate of lead sulfate now obtained will be free from
bismuth. The treatment of the first sulfate precipitate with hydrochloric
acid will dissolve a part or all of the lead sulfate.

4. Dissolve the precipitate of lead sulfate by pouring small portions of

hot 3-normal ammonium acetate solution through the filter. Do not use over
25 cc. of the acetate solution and wash once with hot water after each addition
of the acetate. To thesolution thus obtained add a few drops of potassium
chromate solution and 2 to 5 cc. of 6-normal acetic acid. A yellow precipitate
of lead chromate, insoluble in acetic acid, shows the presence of lead. Bismuth
chromate is readily soluble in acetic acid.

5. Neutralize the filtrate obtained in (3) with ammonia, using a liberal
excess. This serves to precipitate bismuth hydroxide and form soluble com-
plex cations with copper and cadmium. If the original hydrogen sulfide pre-
cipitate was not washed thoroughly, a precipitate of ferric hydroxide or of

aluminium hydroxide may be obtained at this point. Both of these pre-.

cipitates have a different appearance tiian bismuth hydroxide and do not give
the confirmatory test for bismuth. Examine the filtrate by (7).

6. Pour a little 6-normal hydrochloric acid through the filter containing
the bismuth hydroxide and evaporate till only a few drops of liquid remain,
or a moist residue. Add 1 or 2 cc. of water and pour into 100 cc. of hot water.
After standing two or three minutes, filter off any BiOCl and wash once with
cold water. Prepare a fresh solution of sodium stannite by taking a few drops
of stannous chloride solution, diluting with 5 cc. of water and adding sodium
hydroxide solution, a few drops at a time until the Sn(OH), which first forms
redissolves. Pour this solution through the filter containing the BiOCl A
black residue of bismuth is obtained when this element is present (cf. p. 213).

7. If the ammoniacal filtrate from (5) is blue, copper is present. The only
other element which is likely to be confused with copper at this point is
nickel, which also forms a blue solution with ammonia. The color of the
ammoniacal nickel solution is very much paler than that of the corresponding

™

ANALYSIS OF GROUP II 275

copper solution, and it is rare that enough nickel is left with the hydrogen
sulfide precipitate, due to incomplete washing, to cause trouble. In cases
of doubt, acidify a little of the ammoniacal solution with acetic acid and add
_ one drop of potassium ferrocyanide solution. <A red precipitate of Cue[Fe(CN)c],
is formed if copper ts present. The test is much more sensitive than the blue
test with ammonia; nickel, under the same conditions, gives a less character-
istic green precipitate.

(8) If the ammoniacal solution is blue, add potassium cyanide solution
until it is colorless, or only a few drops if the solution is already colorless. Pass
hydrogen sulfide into the solution for half a minute. A yellow precipitate of
cadmium sulfide is formed if cadmium is present. A red precipitate of
(CSNH:)2 may be formed if too much H,S is added (p. 219).

Sometimes a small black precipitate is obtained in the cadmium test which
is most likely due to a little mercury or lead that was not properly removed
from the solution. In such cases, filter off the precipitate, wash it with water
until free from cyanide and boil gently for five to ten minutes in a covered
dish with about 15 cc. of 1.2-normal sulfuric acid (6-normal acid diluted with
_ four volumes of water). This serves to dissolve cadmium sulfide, but will not
dissolve copper or mercury ‘sulfides and should convert lead sulfide into lead
sulfate. Filter, dilute with three times as much water and saturate with
hydrogen sulfide. A yellow precipitate of cadmium sulfide will now be
obtained if cadmium is present.

TABLE IX.—ANALYSIS OF THE ARSENIC-TIN GROUP

Precipitate from Table VII: As.S;, SboS;, SnS2. Warm with 12-normal HCl. (1)

Residue: As.S;. Dissolve in 6-
normal HCl and _ KClIOs.

Solution: Sbt++, Snt++++, Evaporate to small
volume, pour upon clean platinum foil, and place

Evaporate, dilute, neutralize
with NH,OH and add MgCl. -
NH.Cl solution. A white pre-
cipitate of MgNH,AsO, indi-
cates As. Dissolve in 6-nor-
mal HCl, and treat with H2S.
Yellow precipitate of AsS5

a clean piece of zinc in the solution; a black spot
on the platinum indicates Sb. When the evolu-
tion of hydrogen has ceased, remove the zine and,
if any tin deposit adheres to the platinum, rub it
off and dissolve it, with the remaining zinc, in a
small test-tube in one or two drops of concentrated
hydrochloric acid. Dilute with water and add a

or AsS3 shows As. (2) few drops of mercuric chloride solution. A white

or gray precipitate shows Sn is present. (3)

PROCEDURE

1. Transfer the ‘sulfides of arsenic, antimony and tin, which are mixed
with considerable free sulfur, to a porcelain dish and digest for about fifteen
minutes with 12-normal hydrochloric acid. Heat gently, but do not boil
hard. This serves to dissolve antimony pentasulfide (forming antimony tri-
chloride) and stannic sulfide (forming stannic chloride) but does not dissolve
much if any arsenic pentasulfide. Dilute with a little water and filter. Treat
the solution by (8).

2. Heat the residue of arsenic sulfide with 15 ec. of 6-normal hydrochloric
acid and add powdered potassium chlorate from time to time in small portions.

276 HYDROGEN SULFIDE GROUP

When all the yellow sulfide has dissolved, dilute somewhat and filter off the ‘3

‘

residual sulfur. Evaporate the solution to about 2 cc., add 6-normal ammonia
in liberal excess and a few drops of magnesium-ammonium chloride reagent.

If no precipitate forms at once, stir vigorously and rub the sides of the glass
vessel with the stirring rod and let the solution stand at least an hour. Ifa _
precipitate of magnesium ammonium arsenate, MgNH,AsQ,, forms, pour —
off the liquid through a filter and dissolve the precipitate in 10 ce. of 6-normal —

hydrochloric acid. Saturate the solution with hydrogen sulfide, heat to

boiling and again introduce hydrogen sulfide for at least five minutes. A yellow

precipitate of arsenic sulfide will form if as much as 1 mg. of arsenic is present.
3. Evaporate the solution of SbCl; and SnCl,, obtained in (1), to a very

small volume and pour a few drops of the concentrated solution upon a piece ©

of clean platinum foil. Place a small piece of bright zinc in the liquid so that
it comes in contact with both solution and platinum. After a few seconds,

take away the zinc and notice whether a coal black spot of antimony has formed —

upon the platinum. Mercury and copper if present will also be precipitated
_ upon the platinum, but the mercury deposit is gray and the copper deposit

is red. ‘The antimony spot is insoluble in sodium hypochlorite solution

(cf. pp. 235, 249).

Replace the zinc upon the platinum foil and allow the action to continue
until there is no more evolution of hydrogen. Then rinse carefully with dis-
tilled water, taking care not to disturb the contact of the zinc with the platinum.
Take away the zinc and carefully remove any tin that may adhere to the plati-
num. Dissolve the mixture of tin and zinc in one or two drops of concentrated

hydrochloric acid. Dilute the solution in a very small test-tube with a little —

water and add a few drops of mercuric chloride solution. A white precipitate
of mercurous chloride which turns gray shows the presence of tin.

SUPPLEMENTARY PROCEDURES
Analysis of Arsenic—Tin Group

A. If the sulfide precipitate, obtained by acidifying the ammonium
polysulfide solution, consists chiefly of arsenic pentasulfide, it is best to
dissolve the arsenic sulfide first by means of ammonium carbonate solution
(p. 230). Then the residue of Sb.S;, SnS. and § can be examined for antimony
and tin.

To detect the arsenic in the ammonium carbonate solution, add hydro-

chloric acid, which reprecipitates arsenic pentasulfide. Dissolve the sulfide
as indicated in Table X and confirm the arsenic test as described there.

B. An excellent method for separating antimony and tin is based upon the
fact that antimony trisulfide is less soluble than the sulfides of tin. To accom-
plish the separation, proceed as follows:

Dissolve the sulfides of antimony and tin as described above in exactly —

10 ec. of 12-normal hydrochloric acid; when there is no further evolution of
hydrogen sulfide gas, dilute with 3 cc. of water and filter with the aid of suc-
tion. Dilute the filtrate to exactly 55 cc. and transfer the solution to a small
flask. Place the flask in a beaker of boiling water and introduce hydrogen
sulfide into the solution in the flask for ten minutes but no longer. Filter
and wash the precipitate with hot water.

An orange precipitate of antimony trisulfide will be formed when only

a eS — Fy eh

i es 4 A * *
ae Se a ae

ANALYSIS OF GROUP II | 277

1 mg. of Sb is present. Dissolve the precipitate in a little 12-normal hydro-
_ chlorie acid, evaporate nearly to dryness, dilute and test with platinum and
zine as described above.

The filtrate from the Sb.S; can be used for the tin test. Instead of reduc-
ing the tin with zinc, metallic lead can be used (p. 254). Evaporate the solu-
tion to about 3 cc., but not to dryness on account of the danger of losing
_ stannic chloride by volatilization, add 10 cc. of water, 10 gms. of granulated
lead and heat gently for ten minutes. Pour the hot solution through the
filter into 10 cc. of 0.2-normal mercuric chloride. Filter off any mercurous
chloride that may be precipitated, dissolve any adhering lead chloride by wash-
ing with hot water and, if not already blackened, add ammonia to the precipi-
tate (cf. p. 201).

GROUP I. HYDROCHLORIC ACID GROUP

To this group belong silver, mercury (in mercurous compounds),
lead, thallium, and (under some circumstances) tungsten.

SILVER, Ag. At. Wt. 107.88
Sp. Gr.=10.5. M. Pt.=961° C.

Occurrence.—Silver occurs both native and combined (chiefly with
sulfur, arsenic, and antimony).

Of the silver-bearing minerals the following may be mentioned:
horn silver, AgCl; argentite, AgeS; pyrargyrite, AgsSbS3; and prous-
tite, AgsAsSz3.

Silver is also found with tetrahedrite, and with galena.

Metallic silver is of a pure white color. It absorbs oxygen in the
molten state, which it gives up (with tiny explosions) on cooling.

The proper solvent for silver is nitric acid. It is insoluble in
dilute hydrochloric and sulfuric acids, but dissolves readily in boiling
sulfuric acid, with evolution of sulfur dioxide:

Ago+2H2SO4 = 2H20+S802+AgeSOz.

The solubility of silver in concentrated sulfuric acid is utilized
in separating silver from gold and platinum in alloys.

Silver forms three oxides: silver suboxide, Ag4sO; silver oxide,
AgeO; silver peroxide, AgeOo. '

Of these oxides, AgeO alone is a basic anhydride; only one series
of salts is known.

Silver peroxide is formed at the anode during the electrolysis of a con-
centrated, aqueous solution of silver nitrate. The grayish-black crystalline
powder has the composition: Ag,O,=Ag,0-Ag,0;.*

Silver oxide, AgeO, is a brownish-black powder, which on being
heated to 300° is completely decomposed into metal and oxygen.

Most of the silver salts are colorless; the following, however, are col-
ored: the bromide (pale yellow), the iodide (yellow), the sulfide (black),

* BRAUMER and Kuzma, Ber., 1907, 3371. It may be regarded as the silver salt
of the hypothetical argentic acid, HAgOs; ; its symbol could then be written Biri
analogous to the aurates, e.g., K[AuO,].

278

SILVER | 279

the phosphate (yellow), the arsenite (yellow), the arsenate (brown),
the ferricyanide (orange), and the chromate (reddish-brown). Most
of the salts are insoluble in water, and are blackened on exposure to
the light. The nitrete, chlorate, perchlorate, fluoride, nitrite, sulfate,
and acetate are soluble in hot water.

REACTIONS IN THE WET WAY

1. Potassium and Sodium Hydroxides precipitate brown silver
pxide,
2Age*+20H- — H20+Age0,

insoluble in an excess of the precipitant, but readily soluble in nitric
acid and in ammonia. If the solution in ammonia is allowed to stand,
black detonating silver is deposited, [AgNH3]20.

2. Ammonia.—If a neutral solution of a silver salt is cautiously
treated with ammonia, the first drops produce a white precipitate,
which changes quickly to the brown oxide, Ag2O. The greater part
of the silver, however, remains in solution as complex silver-ammonia
salt; even the oxide is dissolved by an excess of ammonia:

Ag20+4NH3+H20 — 2/Ag(NH3)2]*+20H-.

3. Sodium Carbonate precipitates white silver carbonate, which
becomes yellow on being boiled, being slightly decomposed into oxide,
with loss of carbon dioxide:

2Ag*+CO37 => AgeCOs;
AgeCOz ae Age2O+ COeg T °

4. Ammonium Carbonate produces the same precipitate, but
it is soluble in an excess of this reagent.

5. Sodium Phosphate throws down in neutral silver solutions
a yellow precipitate of silver phosphate:

3Agt+ 2HPOs — HePO4-+AgsPOz.

Silver phosphate is easily soluble in nitric acid and in ammonia.
The solution of the phosphate in ammonia is due to the formation
of complex silver-ammonia ions.

AgsP04+6NH3 — 3[Ag(NHs3)2]*+PO4=.

By neutralizing the ammoniacal solution with nitric acid, or the
nitric acid solution with ammonia, the silver phosphate is repre-
cipitated.

280 HYDROCHLORIC ACID GROUP

6. Hydrochloric Acid and Soluble Chlorides precipitate from s
neutral and acid solutions white, curdy silver chloride: ‘aa

Agt+Cl-~— AgCl.

Silver chloride is appreciably soluble in pure water (ef. p. 21), particu-
larly on boiling, but it is quite insoluble if an excess of silver nitrate or of
hydrochloric acid is present (cf. p. 20).

Silver chloride dissolves to a considerable extent in a large excess of hydro-
chloric acid or of alkali chloride, due to the formation of complex anions,
[AgCl;]-, but it is much less soluble in dilute nitric acid.

It is very soluble in ammonia,

AgCl+2NH; — [Ag(NH;).]*CI-,
but is reprecipitated on addition of nitric acid to this solution:
[Ag(NH;).]+2Ht — 2NH,*+AgCl.
Silver chloride is also readily soluble in potassium cyanide solution,
AgC1+2CN~ — [Ag(CN).]-+CI,
and if such a solution is acidified, silver cyanide will be precipitated;
Ag(CN),-+H*t — HCN+AgCN;

In the absence of acid, silver chloride is also readily soluble in a solution
of sodium thiosulfate:

2AgCl+ 38,0;> => [Age (S20 3) 3} si of 7 om

On boiling this solution, silver sulfide is precipitated slowly if considerable _
Na.$.O; is present, but more quickly upon diluting. |

It is evident that the solubility in all the above cases is due to the formation :
of complex ions containing silver (cf. p. 24).

Silver chloride is slowly attacked by boiling, concentrated sulfuric acid,
with evolution of hydrochlori¢ acid, and the formation of crystalline silver
sulfate, insoluble in sulfuric acid.

By boiling with caustic soda or caustic potash solution, silver chloride is
only partially decomposed; in the cold it is unaffected. Sodium carbonate
solution does not affect it; but by fusing with sodium carbonate it is com-
pletely decomposed:

4AgC]1+2Na,CO; =4NaCl+2CO, T +0, fT +4Ag.

By fusing silver chloride itself a yellow liquid is obtained, which on siti
solidifies to a horny mass.

7. Potassium Iodide precipitates yellow, curdy silver iodide,
practically insoluble in ammonia, but easily soluble in potassium
cyanide and sodium thiosulfate solutions. |

The [Ag(NH,).]+ cation furnishes more Ag* cations than silver iodide

in contact with water, but the [Ag(CN).]~ and [Ag,(8.0;),]>~ are much more
stable complexes and furnish less Agt cations than silver iodide does.

SILVER 281

8. Ferrous Sulfate precipitates gray metallic silver from boiling
solutions:

Agt+Fet*— Fettt+Ag.

Frequently a basic ferric salt is precipitated at the same time,
particularly from very dilute solutions. :

9. Zinc, having a greater solution pressure than silver, precipi-
tates the latter from neutral solutions. Similarly, if insoluble silver
chloride is covered with dilute sulfuric acid and a piece of zine placed
in contact with the chloride, the latter will be reduced to metal.

2AgC1+Zn=Zn*++2Cl-+2Ag.

10. Hydrogen Sulfide precipitates from neutral, ammoniacal, and
acid solutions black silver sulfide,

2Ag+H2S > 2H*+Ag.S,

insoluble in ammonia, alkali sulfides, and dilute potassium cyanide
solution. Silver sulfide is so insoluble (cf. p. 21) that enough Agt
ions are present in aqueous solutions containing [Ag(NH3)e]* or even
[Ag(CN)2]— to exceed the solubility product of AgeS when in contact
with slightly ionized hydrogen sulfide. Silver sulfide is, however,
perceptibly soluble in a concentrated solution of potassium cyanide,
and easily soluble in hot dilute nitric acid (cf. p. 181).

11. Potassium Chromate precipitates brownish-red silver chro-
mate, soluble in ammonia and in nitric acid:

2Agt+ CrO4— FRB AgoCrO,.

12. Potassium Dichromate precipitates reddish-brown silver di-
chromate soluble in ammonia and in nitric acid:

2Ag*+Cre07- > AgeCr2O7.

REACTIONS IN THE DRY WAY

Fused with soda on charcoal, all silver compounds yield a white,
malleable, metallic button without incrustation (difference from
lead and tin), readily soluble in nitric acid (difference from tin). The
solution is not precipitated by very dilute sulfuric acid, but is imme-
diately precipitated by hydrochloric acid (difference from lead).

The reactions of lead and mercurous compounds have already been
described (pp. 201, 207).

282 HYDROCHLORIC ACID GROUP

Analysis of Group I

Silver, lead and mercurous chlorides are difficultly soluble in water. +Accord-
ing to the table on page 21, 0.088 mg. Hg.Cl, 0.015 mg. AgCl and about 1
gm. PbCl, will dissolve in 100 cc. of water. From these values, it is evident
that the precipitation of mercurous and silver cations will be nearly complete
from 100 ce. of solution and that, although lead chloride is much less soluble
in a solution containing an excess of chlorine ions than it is in water, lead
will never be precipitated completely as chloride from aqueous solutions. By
the addition of a large excess of alcohol, however, the precipitation of lead
chloride can be made nearly complete, but the addition of alcohol at this stage
of the analysis is rarely permissible except in the analysis of certain alloys
or compounds rich in lead.

Lead chloride is characterized by the fact that its chloride is much more
soluble in hot water than in cold. In dissolving lead compounds in a small
quantity of hydrochloric acid, the chloride often separates out as the solu-
tion cools, and in such cases it is easier to remove the lead chloride at this stage

of the analysis than to precipitate it as sulfide, dissolve in nitric acid and _

reprecipitate as sulfate.

As group precipitant either hydrochloric acid or ammonium chloride can
be used. Any other soluble chloride would accomplish the same end but would
interfere more with the subsequent analysis. If the solution is originally
alkaline, the addition of hydrochloric acid may cause the formation of some
other precipitate. Thus an alkaline cyanide solution, a solution of thio salt
in ammonium sulfide and an alkaline silicate solution will usually give precipi-
tates upon the addition of any acid. Such precipitates often dissolve upon
the addition of a little more acid and cause no trouble, but sometimes it is
best to treat the precipitate as an insoluble substance by the methods
described in Part IV.

Bismuth, antimony and stannic chlorides on being diluted with water
are changed into insoluble basic salts. These precipitates may be filtered off
and dissolved by repeatedly passing a little 2-normal hydrochloric acid through
the filter. Then, as the solution thus obtained is likely to contain lead chloride,
it is best to evaporate to about 1 cc., dilute with 25 cc. of water and precipitate
with hydrogen sulfide without paying any attention to the formation of a basic
salt upon diluting: such basic salts are changed by hydrogen sulfide to less
soluble sulfides.

TABLE X.—ANALYSIS OF THE SILVER GROUP

Solution may contain all the metals. Add 6-normal HCl, filter and examine the
filtrate for succeeding groups. Treat with hot water. (1)

Residue: AgCl, Hg:Ch. Pour ammonia | Solution: Pbt++. Test for lead with

the filter. (3) HSO,. Filter off PbSOsand treat the
precipitate with hot NH4C2H;O2 solu-
Residue: Solution: [Ag(NH;).]. tion. Add KsCrO.; a yellow proces

Hg(NH.)Cl+Hg.| Add HNO;: white tate of PbCrOs shows presence of Pb.
precipitate shows the (2)
presence of Ag. (4)

SILVER 283

The analysis of the first group of metals is based upon the solubility of
lead chloride in hot water, the solubility of silver chloride in ammonia, and the
blackening of mercurous chloride by ammonia.

PROCEDURE

1. To the cold, concentrated solution add 6-normal hydrochloric acid
and, if a precipitate is formed, continue adding hydrochloric acid, drop by
drop until no further precipitation takes place. Filter through a small filter
(ef. p. 58) and wash the precipitate four times with a little cold water, blow-
ing a fine stream from the wash bottle around the upper edge of the filter and
waiting each time until the filter has drained before adding a fresh portion of
water. Do not use more than 5 cc. of water each time. Take the entire
filtrate for the analysis of Groups II, III, IV and V. Pour a little hot water
through the filter. Test the residue by (3).

2. Test the first 5 cc. portion of hot water that runs through the filter
for lead by adding a few drops of concentrated sulfuric acid. If a precipitate
forms, filter, wash once with cold water and then pour 10 cc. of hot 3-normal
ammonium acetate solution through the filter. Add a few drops of potassium
chromate solution to the solution and 3 cc. of 6-normal acetic acid. A yellow
precipitate of lead chromate should form if lead is present. Continue washing
the chlorides with hot water until 5 ec. of the washings will give no test for
lead with potassium chromate solution.

(3) Pour 5 ec. of 6-normal ammonia around the upper edge of the filter,
catching it, as it runs through the funnel, in a test-tube. Acidify the filtrate
with nitric acid. A white, curdy precipitate of silver chloride is formed when
silver is present. The treatment with ammonia causes any mercurous chloride
to turn black on the filter (cf. p. 201).

If considerable mercurous chloride and very little silver chloride is present,
the treatment with ammonia may fail to dissolve silver chloride. In such a
case, wash the black residue and pour repeatedly a mixture of 3 cc. 6-normal
hydrochloric acid and 10 cc. of saturated bromine water through the filter.
This serves to convert the mercurous chloride to soluble mercuric salt and
leaves the silver on the filter as silver chloride. Wash the filter with hot water
and again pour 5 cc. of ammonium hydroxide through it and test the am-
moniacal solution for silver with nitric acid.

PART IIl es

REACTIONS OF THE ACID CONSTITUENTS OR ANIONS

DIVISION OF THE ACIDS INTO GROUPS

The classification of the acids which is given here was first pub-'
lished by R. Bunsen in 1878 (in manuscript form) for the use of his
students; it was adopted (with his consent) by V. Meyer and F. P.
Treadwell in their ‘‘ Tabellen zur qualitativen Analyse.’”’ It is based
upon the different solubilities of the barium and silver salts.

Group I

Acids whose silver salts are insoluble in water and in nitric acid,
but whose barium salts are soluble in water.

To this group belong hydrochloric, hydrobromic, hydriodic, ferro-
cyanic, ferricyanic, cobalticyanic, thiocyanic, and hypochlorous acids.

Group II

Acids whose silver salts are soluble in nitric acid, but are insoluble,
or difficultly soluble, in water, and whose barium salts are soluble in water.
To this group ‘belong hydrosulfuric, -hydroselenic, hydro
nitrous, acetic, cyanic, and hypophosphorous acids.

Group III

Acids whose silver salts are white and soluble in nitric acid, but whose
barium salts are difficultly soluble or insoluble in water, but soluble in
nitric acid. ;

To this group belong sulfurous, selenous, tellurous, phosphorous,
carbonic, oxalic, iodic, boric, molybdic (selenic and telluric), tartaric,
citric, meta- and pyrophosphoric acids.

Group IV

Acids whose silver salts are colored and soluble in nitric acid, but
whose barium salts are insoluble in water and soluble in nitric acid.
To this group belong phosphoric, arsenic, arsenious, vanadic,
thiosulfuric, chromic, and periodic acids.
284

HYDROCHLORIC ACID 285

Group V

Acids whose silver and barium salts are soluble in water.
To this group belong nitric, chloric, perchloric, persulfuric and the

manganic acids.
Group VI

Acids whose silver salts are soluble in water, but whose barium salts
are insoluble in nitric acid.
To this group belong sulfuric, hydrofluoric, and fluosilicic acids.

Group VII

Non-volatile acids, which form soluble salts only with the alkalies.
To this group belong silicic, tungstic, titanic, niobic, tantalic,
and zirconic acids.
GROUP I

Silver Nitrate produces a precipitate insoluble in nitric acid.
Barium Chloride causes no precipitation.

HYDROCHLORIC ACID, HCl

Occurrence.—Hydrochloric acid is found free in nature, but in
small quantities (for example, in the exhalations of active volcanoes) ;
its salts, however, are exceedingly common, especially those with the
alkalies. (See these.) .

Preparation.—Hydrochloric acid is set free by the action of con-
centrated sulfuric acid upon a chloride. Ordinary rock salt is usually
used, it being the cheapest chloride.

If sulfuric acid is allowed to act upon ordinary salt, a consider-
able evolution of hydrochloric acid takes place even in the cold, with
the formation of sodium bisulfate:.

NaCl+H2S04 = NaHS04+HClf¢ .

And, on warming, the sodium bisulfate reacts with more sodium
chloride:
NaHS0O4 +NaCl = NaeS04+ HCl.

When only one hydrogen of a dibasic acid is replaced by metal,
the resulting salt is usually designated by placing the prefix bi before
the name of the acid; sodium bicarbonate, NaHCO3; sodium bisul-
fate, NaHSOa, etc. There is twice as much acid per equivalent of base
a3 in the case of the normal salt.

Hydrochloric acid may also be prepared by dropping concentrated

286 REACTIONS OF THE ACID CONSTITUENTS

sulfuric acid into 12-normal hydrochloric acid. It is formed by the
action of water on many acid chlorides:

PCl3+3H20 = H3P03+3HCI.

Properties.—Hydrochloric acid is a colorless gas, with a suffocating
odor, which forms dense clouds in moist air. It is readily soluble in
water (one volume of water dissolves, at 18° C., 451 volumes of hydro-
chloric acid). The specific gravity of the saturated, aqueous solu-
tion is 1.21, and 100 cc. of this solution contain 50.7 gms. of hydro-
chloric acid gas. The concentrated hydrochloric acid of the laboratory
is about 12-normal, and contains 36 to 38 per cent of the gas by
weight. As regards its behavior on evaporation, see page 57. The
aqueous solution of hydrochloric acid is one of our strongest acids.
In dilute solution it is almost entirely ionized,

HCl = H*+Cr,

and such a solution is a good conductor of electricity.
The behavior of hydrochloric acid on oxidation is extremely character-
istic; water is formed and chlorine 1s set free:

2HC1+0 = H20+Cle.

Tnis oxidation will not take place on exposure to atmospheric,
or even pure, oxygen, but only by strong oxidizing agents such as:

The peroxides of the heavy metals, all nitrates, nitrites, chlorates,
hypochlorites, chromates, selenates, and tellurates.

The peroxides of the light metals do not yield chlorine, but hydrogen
peroxide:
BaO.+2HCl =BaCh.+H:202.

1. Oxidation of Hydrochloric Acid by Means of Peroxides

(a) By manganese dioxide:
4HCl+ Mn0O2=2H20+ MnCle+Cle fT.

The manganese is reduced from the quadrivalent to the bivalent condition
and the negatively charged chlorine atom becomes neutral chlorine gas. Thus
one atom of manganese oxidizes two atoms of chlorine.

Some other acid, preferably sulfuric, can be used to combine with
the manganese and then the yield of chlorine from a given quantity
of hydrochloric acid will be twice as large:

MnO2+2HCl+H2804= Mn804+2H20+Cle f.

rl 4

— ae PF ip

p™

HYDROCHLORIC ACID 287

(b) By lead peroxide:
4HCI+PbO2 = 2H20+PbCle+Cle fF .
(c) By chromium trioxide:
2CrO3+12HCl = 6H20+2CrCl3+3Cle.

2. Oxidation of Hydrochloric Acid by Nitric Acid, etc.

When concentrated nitric acid acts upon concentrated hydro-
chloric acid, the products of the reaction are water, nitrosyl chloride

and free chlorine:
HN0O3+3HCl=2H20+NOCl Tf +Cle ft.

In this reaction the valence of the nitrogen has been reduced from
five to three, and two atoms of chlorine have been oxidized to form
an electrically neutral chlorine molecule.

A mixture of one molecule of nitric acid with three of hydrochloric
acid is known as

Aqua Regia.—The acids are usually mixed, not desta ae to their
weights, but according to their volumes.

Aqua regia is, therefore, chlorine water, with the distinction that
the chlorine exists in the nascent state; which explains why aqua
regia is a much more energetic reagent than ordinary chlorine water.
Nitrous acid, chloric acid, hypochlorous acid, selenic and telluric
acids all react similarly with hydrochloric acid.

Hydrochloric acid is monobasic; its salts are called chlorides.

Solubility of Chlorides

Nearly all chlorides are soluble in water, but the following are
insoluble: mercurous chloride, HgeCle; silver chloride, AgCl; cuprous
chloride, CueCle; lead chloride, PbCle; thallium chloride, TIC;
aurous chloride, AuCl; platinous chloride, PtCle; bismuth oxychlor-
ide, BiOCl; antimony oxychloride, SbOCl; mercuric oxychloride,
HgeCl.O.

All these chlorides which are insoluble in water are more soluble
in strong hydrochloric acid.

Aurous chloride and platinous chloride dissolve readily in aqua
regia, as a result of being oxidized, but silver chloride is not very solu-
ble even in aqua regia.

By boiling the insoluble chlorides with a concentrated solution
of sodium carbonate, all, with the exception of silver chloride, are
readily decomposed, e.g.:

Hg2Cle +NaeCO3= = 2NaCl+CO2 q +Hg.0.

‘

288 REACTIONS OF THE ACID CONSTITUENTS _

By filtration.a chloride solution is obtained which is free from
heavy metal.
By fusing with sodium carbonate, even silver chloride is decom-
posed,
4AgC1+2Na2CO3 =4NaCl+2CC2 T +02 T +4Ag,

and silver chloride may also be decomposed by nascent hydrogen
(ef. p. 281).

The deliquescent chlorides (lithium, calcium, and_ strontium)
are all soluble in absolute alcohol and in amyl alcohol.

The chlorides of potassium, sodium, and barium are quite insoluble
in concentrated hydrochloric acid; they can, therefore, be easily

separated from the remaining chlorides which are soluble in water

by saturating the solution with hydrochloric acid gas.
Almost all chlorides are insoluble in ether, with the exception of
mercuric, stannous, stannic, auric, and ferric chlorides,

REACTIONS OF CHLORIDES IN THE WET WAY

A neutral solution of an alkali chloride should be used for these
reactions. All chlorides except those of mercury and cadmium are
almost completely ionized in aqueous solution, so that it is a matter
of indifference which chloride we take for the following reactions,
provided there is no independent reaction taking place due to the
presence of the cation.

1. Dilute Sulfuric Acid a 10) produces no reaction, even on
warming. |

2. Concentrated Sulfuric Acid decomposes the solid chloride
almost completely in the cold, completely on warming. Sulfate
and colorless hydrochloric acid gas result from this reaction; and
the latter is easily recognized by its odor, by the clouds which it forms
in moist air or with ammonia vapors (obtained by holding a glass rod
wet with ammonia near the test-tube), and by its turning moist blue
litmus-paper red. Water is not made turbid by hydrochloric acid
(difference from fluosilicic acid).

Silver chloride and mercurous chloride are decomposed with difficulty by
sulfuric acid, the latter with evolution of sulfur dioxide; the mercurous sul-

fate (which is at first formed) is oxidized (at the expense of the oxygen of the
sulfuric acid) to mercuric sulfate:

(a) Hg.Cl, -{- H.SO, = Hg.SO,+ 2HCl >
(b) Hg.S0O,+ 2H.SO, = 2HgSO.+ 2H.0 + SO, T ;

If a chloride and an oxidizing agent are heated with concentrated sulfuric
acid, free chlorine is evolved (cf. p. 286).

HYDROCHLORIC ACID 289

3. Phosphoric Acid heated with a chloride similarly causes the
evolution of hydrochloric acid gas. If an oxidizing agent is present at
the same time, chlorine gas is evolved.

4. Silver Nitrate produces a white, curdy precipitate of silver
chloride, .
| Cl-+Ag* — AgCl,

insoluble in nitric acid, soluble in ammonia, potassium cyanide, and
sodium thiosulfate solutions. (See Silver.)

From a solution of silver chloride in ammonia, acids reprecipitate silver
chloride. From a solution in potassium cyanide, acids precipitate silver cyanide.
If it is desired to test a solution of ferrous sulfate for the presence of a small
amount of chloride, it must be strongly acidified with nitric acid, as otherwise
a precipitate of metallic silver will be obtained, which may cause confusion
(ef. p. 281). The best way to test the solution of ferrous sulfate for hydro-
chloric acid is to add sodium carbonate solution until alkaline, boil, and filter.
In the filtrate the acids originally present are now in the form of their sodium
salts, in the presence of an excess of sodium carbonate; and the latter should
be neutralized with nitric acid before the silver nitrate is added.

The detection of chlorine when present in the form of chloride of a heavy
metal is accomplished in a similar manner; and, with the exception of silver
chloride, any insoluble chloride may be decomposed in the same way, by
boiling with sodium carbonate solution.

In order to detect the presence of chlorine in silver chloride, treat it with
dilute sulfuric acid and zine (p. 281), after a short time pour off the solution
from the deposited silver and test it with silver nitrate.

Or, fuse the silver chloride with sodium carbonate, extract the melt with
water, filter, acidify with nitric acid, and test with silver nitrate.

5. Potassium Dichromate and Sulfuric Acid.—If a dry chloride
is mixed with potassium dichromate, concentrated sulfuric acid is
added, and the mixture heated in a small retort, brownish vapors are
given off which condense, in the receiver, to a brown liquid (chromyl

chloride, CrO2Cle):
KeCre07+4NaCl+ 3H2804 = 3H20 + 2Na28041+ Ko804+ 2CrO2Cle Ty.

Chromy] chloride is an acid chloride, and: is, therefore, decomposed
by water into chromic and hydrochloric acids:

CrO2Cle+2H20 = HeCrO4+2HCl.

On adding caustic soda or potash, an alkali chloride and a yellow
alkali chromate are obtained. If the solution is then acidified, some
ether and a little hydrogen peroxide added, and the liquid shaken,
the upper ether layer will be colored blue, showing the presence of

290 REACTIONS OF THE ACID CONSTITUENTS

chromium; and the presence of chromium indicates that a chloride
was originally present (difference from bromide and iodide).
6. Potassium Permanganate oxidizes a chloride even more readily.

On the other hand, hydrobromic and hydriodic acids are oxidized |

much more readily than yi acid.

Behavior of Chlorides on Ignition
The chlorides of the alkalies and alkaline earths melt (without

perceptible decomposition) on being heated in the air. The chlorides —

of the sesquioxides are decomposed, more or less completely, on being
ignited in the air. Thus, ferric chloride is almost quantitatively
decomposed into ferric oxide, with loss of chlorine:

2FeCls+30 = FeeQO3 +3Cle.

In the presence of hydrochloric acid, or ammonium chloride, ferric chloride
may be volatilized completely without any decomposition.

The chlorides of gold and of the platinum metals are readily
decomposed into chlorine and metal:

2AuCls = 2Au+3Cle;
PtCh =Pt+2Ch 1.

The remaining chlorides are mostly volatile, without perceptible
decomposition.

Detection of Chlorine in Non-electrolytes

Besides uniting with metals and with hydrogen, chlorine also

forms compounds with the metalloids; e.g., PCls, PCls, AsCls, AsCls, —

SbCla, SbCl5, CCla, SiCl, ete.

All these compounds, which may be regarded as acid chlorides, —

are decomposed by water with the formation of hydrochloric acid,
which can be tested for in the usual way. The hydrolysis usually

takes place at the ordinary temperatures. Thus, phosphorus tri-—

chloride yields with water phosphorous acid and hydrochloric acid,
PCls+3HOH = H3P03+3HCl,

and phosphorus pentachloride yields phosphoric acid and hydrochloric

acid:
PCl;+4HOH = H3P04+5HCI.

The remaining acid chlorides are decomposed in a similar way by
water at the ordinary temperature, with the exception of carbon

FREE CHLORINE 291

tetrachloride, which is decomposed by water only by heating in a

closed tube:
st CCl4+2H20 =CO02+4HCl.

Chlorine acts upon a great many hydrocarbons, forming substitution-
products which are non-electrolytes, and consequently will not give the
chloride tests; for example, if chloroform, CHCl;, is shaken with a solution
of aqueous silver nitrate, it will not yield a precipitate of silver chloride. In
order to test such compounds for chlorine (as is frequently necessary in the
study of organic compounds), the chlorine must be changed to hydrochloric
acid by one of the following methods:

1. Carius’ Method.—By heating the compound in a sealed glass tube
with concentrated nitric acid, in the presence of silver nitrate, the compound
is completely decomposed; all the chlorine is changed to silver chloride,
which can be filtered off, and, after treatment with zinc and dilute sulfuric
acid, can be tested as above. The precautions to be taken in sealing and
opening the tube are described in Vol. II.

2. By Heating with Lime.—A layer of granular lime (free from chloride)
then a mixture of the substance to be tested and lime, and finally another
layer of lime are placed in a small glass tube, which should be about 25 em.
long and about 1 cm. wide.

By gently tapping the tube, a canal is opened between the upper wall of
the tube and the substance, through which the gases evolved may escape.
The tube is then placed horizontally in a small combustion furnace and heated
(first the front layer of lime, then the back layer, and finally the entire contents
of the tube) to a dull red heat.

By this means the organic substance will be completely decomposed, and
the chlorine will be found combined with the lime in the form of calcium chloride.

After cooling, the contents of the tube should be dissolved in dilute nitric -
acid, the carbon filtered off, and the filtrate tested with silver nitrate for
chlorine ions.

3. Treatment with Sodium.—A small amount of the dry substance to be
tested is placed in a small test-tube, a small piece of sodium (freed from petro-
leum) is added, and the metal covered with another layer of the substance.

_ The tube is then heated in the gas-flame, the decomposition taking place
suddenly with incandescence. The still hot tube is transferred to a small
beaker containing water (which breaks the tube) and sodium chloride dis-
solves with other sodium compounds. The solution is filtered, acidified with
nitric acid, and then tested with silver nitrate for halogens.

FREE CHLORINE

Chlorine, whether produced by the oxidation of hydrochloric acid
or by igniting certain chlorides, is a greenish-yellow gas, with a suffo-
cating odor. It is absorbed by water (one volume of water absorbs
at 10° C. about 2.7 volumes of chlorine gas), forming chlorine water,
a yellowish-green liquid, and a poor conductor of electricity (although
better than pure water, showing that some ions are present). Chlorine

292 REACTIONS OF THE ACID CONSTITUENTS —

decomposes water to a slight extent, forming hydrochloric acid and
hypochlorous acid:

H20+Cle @ 2Ht+Cl-+Clo-.

In this reaction one atom of chlorine is oxidized to a positive valence of one
and another atom of chlorine is reduced to a negative valence of one (cf.
-p. 170). This reaction, as the application of the mass-action law indicates,
is favored by the presence of hydroxyl ions and hindered by the presence
of hydrogen ions. If, besides water, which is only slightly ionized, dilute alkali
is present, the reaction takes place quantitatively from left to right.

Cl,+OH- > Cl-+ClO-+H:,0.

The decomposition of chlorine water takes place slowly in the
dark, but more rapidly in the light, and in the presence of oxidiz-
able substances. The bleaching action of chlorine water depends
upon its oxidizing power. :

Chlorine water ts a strong oxidizing agent.

If a solution of potassium iodide is treated with chlorine water,
iodine is set free, and the solution turns yellow to brown: /

2KI+Cle=2KCI+Io.

If the yellow solution is shaken with carbon disulfide, or chloroform,
the iodine, a more highly colored, violet, non-aqueous solution of iodine
is obtained (cf. p. 17). By the addition of more chlorine water the
solution becomes colorless, owing to the oxidation of the iodine to
colorless iodic acid:

Io+6He0+5Cle = 10HCI+2HI10Os3.

The free iodine can also be detected by the addition of some starch
paste (instead of carbon disulfide, etc.), which is turned blue by iodine.
1. Silver Nitrate gives a white precipitate of silver chloride when
added to chlorine water; this precipitation, however, is not quantita-
tive, for one-sixth of the chlorine is changed into soluble silver chlorate:

3Clo+6Ag*+3H20 = 5AgCl+AgCl03+6H".

On adding a slight excess of sulfurous acid to chlorine water, the chlorine
is readily and completely changed into hydrochloric acid:

Cl, of H.O -+ H.SO; = H.SO,+ 2HCl.

From this solution the chlorine can be precipitated quantitatively by a
silver solution.

Chlorine may also be easily changed into the form of a chloride by the
action of ammonia (cf. p. 165).

3Cl.+8NH; oy 6NH,CI+N. t .

FREE CHLORINE 293

Chlorine can be changed to a chloride by the action of hydrogen peroxide
in the presence of sodium or potassium hydroxide: ‘

Cl.+20H-~ — OCI-+Cl-+H20,
and
ClO-+H,0, — H:0+ClI-+0, fT.

2. Metallic Mercury is attacked by chlorine at the ordinary tem-
perature, forming insoluble mercurous chloride.

Hge a Cle = Hge2Clo.

If therefore chlorine water is shaken with metallic mercury until it no longer
smells of chlorine, a neutral solution is obtained which contains no chlorine.
If hydrochloric acid is present, the solution now reacts acid, and gives a pre-
cipitate with silver nitrate, for metallic mercury is not attacked by hydro-
chloric acid. ‘This reaction is used as a test for hydrochloric acid in the pres-
ence of chlorine.

3. Metallic Zinc also reacts with chlorine water,

Zn+-Cl, > Zntt+2cCr-.

HypocHLorovus AciID, HOCI

Preparation.—A solution of free hypochlorous acid is obtained _

by shaking chlorine water with yellow mercuric oxide until the solution
no longer smells of chlorine:

2Hg0+2Clo+H20 = (HgCl)20+2HOCI.

Brown, insoluble mercuric basic chloride is formed by the reaction,
and the solution contains hypochlorous acid. If the solution is poured
off from the insoluble basic mercuric salt and distilled, a pure solu-
tion of hypochlorous acid will be obtained; which, however, cannot
be kept long in the light, for it decomposes into hydrochloric acid
and oxygen:

2HOCI — 2HCI+02 7.

Hypochlorous acid is a vigorous bleaching agent; litmus and
indigo are quickly decolorized.

The alkali salts of hypochlorous acid (hypochlorites) are obtained —
by neutralizing the acid with dilute sodium or potassium hydroxide;
or more conveniently, by the action of chlorine on a dilute caustic
alkali solution:

Cle +20H-=Cl-+ClO~+H20.

The ammonium salt cannot be prepared except at very low tem-
peratures, because the nitrogen of ammonia is so readily oxidized (ef.
p. 165).

All hypochlorites are readily changed, on warming, into chlorate *
and chloride (cf. p. 170):

3KCIO — 2KC1+KCl0s,

consequently hypochlorites must always be prepared in cold, dilute
solution.
The most important commercial hypochlorite is the so-called

“ chloride of lime,” which is obtained by passing chlorine gas over
lime at the ordinary temperature.

*In the presence of 40 per cent or more of caustic potash the potassium hypo-
chlorite on being heated decomposes into chloride with evolution of oxygen and the
formation of no perchlorate (F. WiNTELER, Z. angew. Chem., 33 (1902), p. 778).

294

HYPOCHLOROUS ACID 295

REACTIONS IN THE WET WAY

___ All hypochlorites are soluble in water, and are decomposed by
acids (carbonic acid even).
1. Hydrochloric Acid is oxidized by hypochlorites with evolution
of chlorine:
ClO-+Cl-+2Ht > H20+Cle.

The reaction is favored by the presence of hydrogen ions. The
reverse reaction is favored by the presence of hydroxyl ions (ef. p.
292).

2. Sulfuric Acid decomposes hypochlorites, setting free hypo-
chlorous acid: |

NaOCl+ H2804 = NaHSO4+HOCI,

_and carbonic acid acts similarly though less energetically.
2Na0Cl+ HeCO3 = NaeCO3+2HOCI.

It is due to the fact that hypochlorites are so readily decomposed
with the formation of chlorine that they act as strong bleaching
agents, indigo solution (a solution of indigo in sulfuric acid) being
immediately decolorized.

Hypochlorites act as oxidizing agents not only in acid solutions,
but also in alkaline solutions at ordinary temperatures (difference
from chlorates), many metallic hydroxides being oxidized by them
to higher hydroxides. Thus, ferrous hydroxide is readily oxidized
to reddish-brown ferric hydroxide:

2Fe(OH)2+ClO” +H20 — Cl" +2Fe(OH)s,

and similarly lead, manganous; -nickelous, and cobaltous hydroxides
are oxidized to brownish-black hydroxides.
Less than 1 mg. of hypochlorite can be detected by boiling with

lead acetate solution:
Pb*+++ClO-+H:20 — PbO2+2H*+Cr-.

If a peroxide is likely to be present, the solution must first be acidified

with acetic acid before applying the test (cf. p. 86.)
3. Iodo-Starch Paper is turned blue by hypochlorites in weakly
alkaline solutions, owing to the separation of iodine:

2KI+Na0Cl+ H20 =2KOH+NaCl+l..
21-+ClO-+H20 @ Cl-+20H +Iz.
The extent to which the reaction takes place depends upon the

296 REACTIONS OF THE ACID CON STITUENTS

concentration of the hydroxyl ion; only in dilute alkali will enough or

iodine be formed to produce the blue coloration with the starch.

4. Metallic Mercury.—If a solution containing free hypochlorous 3

acid is shaken with metallic mercury, brown basic mercuric chloride
is formed, insoluble in water, but soluble in hydrochloric acid:

2He+2HOC!I = (HgCl)20+ H20.
Under these same conditions free chlorine forms, on being shaken
with mercury, white mercurous chloride, which is insoluble in hydro-
chloric acid.

. This property is utilized in detecting hypochlorous acid in the presence
of free chlorine. The mixture is shaken with mercury until a little of the
solution no longer turns iodo-starch paper blue; the liquid is then carefully
poured off, hydrochloric acid is added to the residue, and the mixture is shaken,
when the basic chloride produced by hypochlorous acid goes into solution:

(HgCl),0+2HCl = H,0+2H¢Ch,

while mercurous chloride remains undissolved.

If hydrogen sulfide is passed into the filtered solution, the formation of
mercuric sulfide shows that hypochlorous acid was originally present.

The salts of hypochlorous acid behave differently toward mercury than
is the case with the free acid; they form insoluble mercuric oxide and a soluble
chloride:

Hg+Na0Cl =HgO+NaCl.

5. Silver Nitrate causes in solutions of hypochlorites an incom-

plete precipitation of silver chloride. One-third of the chlorine remains ~

in solution in the form of chlorate:
3ClO~-+2Ag* — C1037 +2AgCl.

Hypochlorous acid is distinguished from chlorine by its behavior
toward mercury; from hydrochloric acid by its oxidizing action;
and from chloric acid by its being partly precipitated by silver
nitrate, and by its oxidizing action in alkaline solutions. .

HYDROBROMIC ACID, HBr.

Occurrence.—Bromine is always found in nature in company with
chlorine; thus salts of hydrobromic acid are found in the ocean and
in many. mineral waters.

Preparation.—Hydrobromic acid is formed by the action of sul-
furic acid upon a bromide:

2NaBr+ H2S804 = NaeSO4+2HBr.

But the hydrobromic acid obtained by means of this reaction is never
pure, being contaminated with bromine; a part of the hydrobromic
acid is oxidized by the sulfuric acid:

2HBr+ H2SO4 — 2H20+S02 tT +Bre T :

The more concentrated the sulfuric acid used, the larger will be
the yield of bromine.

If dilute sulfuric acid (83H2SO4 : 1H2O) is used, the hydrobromic
acid obtained is nearly free from bromine.

Pure hydrobromic acid may be obtained by: the action of an acid
bromide upon water:

PBr3 +3HOH = Hs3POz3 +3HBr.

Properties. ..Hydrobromie acid (like hydrochloric acid) is a color-
less gas with a suffocating odor, having the property of fuming in
moist air and forming clouds of ammonium bromide with vapors
of ammonia. It is very soluble in water. The concentrated solu-
tion has the sp.gr. 1.78 and contains 82 per cent by weight of hydro-
bromic acid. Hydrobromic acid is less stable than hydrochloric
acid and more stable than hydriodic acid.

While hydrochloric acid can be kept in aqueous solution for an
indefinitely long time, a solution of hydrobromic acid soon turns
brown, owing to the separation of bromine. Hydrobromic acid is

oxidized by atmospheric oxygen:
4HBr+O2 = 2H20+4+2Broe.

Owing to the slight solubility of cuprous bromide, and aided by the
presence of a little free bromine, hydrobromic acid will dissolve copper,
297

298 REACTIONS OF THE ACID CONSTITUENTS

although this element is below hydrogen in the electromotive series
(cf. pp. 38 and 41). :
Hydrobromic acid is readily oxidized with separation of bromine,
by peroxides, nitrates, chromates, etc., provided a concentrated solu-
tion of hydrobromic acid is used.
Hydrobromic acid is a monobasic acid; its salts are called bromides.
The solubility of a bromide is similar to that of the correspond-
ing cbloride.

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid (1:10) evolves no hydrobromic acid
from bromides in the cold, but, on warming, does so from bromides
of the alkalies.

2. Concentrated Sulfuric Acid causes evolution of hydrobromic
acid and bromine from all bromides: |

H2S804+2Br — SO47>+2HBr
and . .
H2$04+2Br — H20+S802 T +Bre T :

The solution is brown, and, on warming, yellowish-brown vapors are
given off (difference from hydrochloric acid); which, as they contain
hydrobromic acid, fume in moist air, have a suffocating odor, and do
not render water turbid (difference from fluosilicie acid).

3. Phosphoric Acid causes the evolution of hydrobromic acid.
If an oxidizing agent is also present, bromine is evolved. |

4. Silver Nitrate produces a curdy, yellowish precipitate of silver
bromide, insoluble in nitric acid, but soluble in ammonia, potassium
cyanide, and sodium thiosulfate. Silver bromide is less soluble than
silver chloride.

If, therefore, silver chloride is digested with potassium bromide,
the former will be changed into silver bromide:

AgCi+KBr=AgBr+ KCl. .

If, however, silver bromide is heated and treated with chlor
gas, it is readily changed into silver chloride:

AgBr+Cl=AgCl+Br.

5. Chlorine Water, on being added to solutions of soluble bro-
mides, sets free bromine, which is soluble in carbon disulfide or chlo-
roform, forming a brown.solution; but it is changed, by an excess of
chlorine water, into yellowish chloride of bromine (BrCl) (difference
from iodine). |

FREE BROMINE 299

6. Potassium Dichromate, in the presence of dilute sulfuric acid,
does not cause separation of bromine from aqueous. solutions of
bromides; on shaking the solution with carbon disulfide, the latter
remains colorless (difference from iodine). |

7. Potassium Dichromate and Concentrated Sulfuric Acid.—
On mixing a solid bromide with solid potassium dichromate, cover-
ing the mixture with concentrated sulfuric acid and distilling, a brown
distillate is obtained (as with a chloride), which, however, consists
of bromine and contains no chromium:

Cr207~+6Br-+14H* — 2Cr* **+7H20+3Bre ¢ .

On adding dilute sodium hydroxide to the distillate, a colorless
(or sometimes a faint yellow) solution is obtained; which, after being
acidified with sulfuric acid, does not give the chromium reaction with
dilute sulfuric acid and hydrogen peroxide, but merely turns brown
owing to the liberation of free bromine (difference from chlorine).

8. Potassium Permanganate oxidizes hydrobromic acid more
readily than it does hydrochloric acid and less readily than it does
hydriodic acid.

9. Nitrous Acid does not cause the separation of bromine from
a dilute bromide solution in the cold (difference from iodine).

Detection of Bromine in Non-electrolytes

The method of procedure is exactly the same as was given in the
case of chlorine in a non-electrolyte (see p. 291).

FREE BROMINE

Free bromine (which may be obtained by the oxidation of hydro-
bromic acid) is a brown liquid at the ordinary temperature, and dis-
solves in water, forming a colored solution. The cold, saturated
solution of bromine contains 2 to 3 per cent of dissolved bromine.
Concentrated hydrochloric acid at the ordinary temperatures dis-
solves much more bromine, the saturated solution containing about
13 per cent of the substance.

Bromine, like chlorine, acts as a strong bleaching agent (oxidiz-
ing the color) and combines directly with metallic mercury, forming
insoluble mercurous bromide.

The detection of hydrobromic acid in the presence of bromine
is accomplished in precisely the same way as was described for the
detection of hydrochloric acid in the presence of free chlorine (cf.

p. 293).

HyYDRIODIC ACID, HI

Occurrence-—Iodine occurs in nature as the iodide and as the
iodate; most frequently as the former, accompanying (in small amounts)
chlorine and bromine; e.g., in the ocean, in mineral waters, etc.

_Preparation.—Hydriodic acid may be obtained pure by the action
of an acid iodide upon water:

PI3+3H20 = H3P03+3HI.

If we attempt to prepare hydriodic acid by the action of sulfuric
acid upon iodides, even from dilute solutions, it is more difficult than
in the case of hydrobromic acid to obtain a pure product, on account
of the strong reducing power of hydriodie acid. The hydriodic acid
thus obtained always contains a large amount of iodine, together
with the reduction products of sulfuric acid, the latter varying in com-
position according to the concentration of the acid employed and of the
iodide solution. Thus, with considerable concentrated sulfuric acid,
sulfur dioxide is obtained:

2Nal+3He2S04 ==> 2NaHS0O4+2H20+S02 T +Io.

But if considerable iodide is present, the sulfuric acid is reduced
to sulfur and finally to hydrogen sulfide.

6KI+7H2804 — 6KHSO4+$+4H20+4+3ly,
8KI+9H2804 — 8KHSO4+ HS 7 +4H20+4]e.

Properties—Hydriodic acid is a colorless gas, with a suffocating
odor; it fumes in moist air, and is readily soluble in water, forming
a strongly fuming liquid of sp.gr. 1.99-2.10. The aqueous solution
of hydriodic acid is even more difficult to keep than a solution of
hydrobromic acid; it soon turns brown, owing to the separation of
iodine:

4HI+Oz2 (air) =2H20+2le.

If hydrogen sulfide is conducted into the brown solution, it -is

decolorized, with separation of sulfur:

I2+H2S=2HI+58.
301

302 REACTIONS OF THE ACID CONSTITUENTS

Owing to the extremely slight ionization of complex mercuric iodide
anions, hydriodic acid is capable of dissolving mercury, which is con-
siderably below hydrogen in the electromotive series (cf. pp, 10 and

37).
Hg+4HI =[Hgl4]H2+ He.

Hydriodice acid (like hydrochloric and hydrobromic acids) is oxi-
dized by peroxides, nitrates, nitrites, chromates, etc., with separa-
tion of iodine; only in this case the oxidation of the hydriodic acid
takes place much more readily, so that a very dilute solution is oxi-
dized by nitrous and chromic acids even in the cold.

Hydriodie acid is a monobasic acid; its salts are called iodides.

The solubilities of the iodine compounds are analogous to the
corresponding bromine and chlorine compounds. The cuprous, siiver,
mercury and palladium salts are much less soluble than the corre-
sponding chlorine or bromine compounds,

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid (1: 10) attacks the iodides of the alkalies
perceptibly, but only on warming.

2. Concentrated Sulphuric Acid reacts in the cold (ef. p. 301).

3. Silver Nitrate produces a yellow, curdy precipitate of silver
iodide, insoluble in nitric acid, and only slightly soluble in ammonia,
but readily soluble in potassium cyanide and sodium thiosulfate.

Treatment with ammonia causes the silver iodide to assume a much lighter
color.

By the action of chlorine gas, silver iodide is readily changed into

silver chloride:
2AeT+Clo=2AgCl+To.

On the other hand, if chloride or bromide of silver is treated with
potassium iodide it will be changed into silver iodide:

AgCl+KI=KCl-+AgI,
AgBr+KI=KBr-+AglI.

This apparently anomalous behavior is easy to explain. As its position in
the electromotive series shows, free chlorine can displace bromine or iodine in
its salts; it hasa greater tendency to be reduced to chloride ions. On the other
hand, silver iodide is far less soluble than silver bromide and the latter is less
soluble than silver chloride, and, for this reason ‘alone, iodine ions will replace
chlorine or bromine ions in the case of the silver salts,

HYDRIODIC ACID 303

4. Lead Salts precipitate yellow lead iodide, soluble in consider-
able hot water and forming a colorless solution which deposits gold-
yellow plates of PbI2 on cooling.

5. Palladous Chloride (it is best to use sodium palladous chloride)
precipitates, from dilute solutions of an iodide, black palladous iodide
(difference from chlorine and bromine).

Nas(P dCl4]+2KI =2NaCl+ 2KCI +PdIo,

which is readily soluble in an excess of potassium iodide.
6. Cupric Salts are reduced by iodides, causing the separation
of a brownish mixture of cuprous iodide and iodine:

Cut te4r —> Cuzl2+ Ie.

If sulfurous acid or sodium thiosulfated is then added to the solution,
a nearly white deposit of cuprous iodide is obtained, owing to the
free iodine being changed to hydriodic acid by the sulfurous acid:

SO3° +H20+I2 —> SO4-+2Ht+2I -;
28203 +Is — 8406 +2I-.

7. Nitrous Acid.—If a dilute solution of an iodide is treated with
nitrous acid, iodine separates out and the solution becomes yellow
or brown in color (difference from chlorides and bromides):

2HNO2+2HI =2NO JT +12+2H20.

This extremely sensitive reaction is best performed as follows:

Prepare a solution of nitrous acid in concentrated sulfuric acid by heating
arsenic trioxide with nitric acid (sp.gr. 1.30-1.35) and -conducting the gases
evolved (NO; and NO) into sulfuric acid (sp.gr. 1.75-1.80):

As.0O;+2HNO; =As,.O;+H,0+ NO, tT +NO tT ;

—OH

This solution of nitrosyl sulfuric acid is sometimes called “ nitrose.” It can
be kept for some time, but water decomposes it into nitrous and sulfuric acids:

80:_pNot HOH ~ =H.S0.+ HNOQ..
Treat the solution to be tested with a few drops of nitrose and shake with

a little carbon disulfide or chloroform; if an iodide is present, iodine is liberated,
which colors the carbon disulfide or chloroform.

304 REACTIONS OF THE ACID CONSTITUENTS

8. Potassium Dichromate, in the presence of dilute sulfuric acid,
causes the separation of iodine from dilute iodide solutions in the
cold; the iodine can be more easily recognized by shaking the solu-
tion with chloroform or carbon disulfide (difference from bromine):

Cro077+61-+14Ht > 2Cr++++7H20+4+3I2:

By heating a mixture of solid iodide and solid potassium dichro-
mate with concentrated sulfuric acid, iodine is set free (according
to the above equation), which distills over, and can be condensed in
the receiver. No chromium is carried over by this reaction (differ-
ence from chlorine). :

9. Potassium Permanganate oxidizes hydriodic acid more readily
than acts upon hydrobromic¢ or hydrochloric acids. The oxidation will
take place in dilute acetic acid solutions containing considerable sodium —
acetate.

10. Mercuric Chloride produces scarlet mercuric iodide, soluble
in an excess of potassium iodide:

Hg*t+2I- — Hels, |
Hgl2+2l" — [Hgly]-.

11. Chlorine Water sets free iodine from iodides,
2KI+Cle=2KCI+Iy,

which colors carbon disulfide reddish-violet, or starch-paste blue.
By adding an excess of chlorine water, the violet color disappears, |
the iodine being oxidized to colorless iodie acid:

Iz+6H20+5Cle = 10HCI+ 2HI0s.

Detection of Iodine in Non-electrolytes

The processes to be employed are the same as those described for
detecting chlorine in non-electrolytes (see p. 291).

FREE IODINE

Free iodine forms scales resembling graphite in appearance, has
a sp.gr. 4.94 at 17°. It melts at 114° (at the same temperature as
sulfur), but begins to volatilize at ordinary temperatures, and is
completely transformed into violet vapors at 200°.

Iodine is only slightly soluble in water (100 parts of water dissolve
0.02 part of iodine) but is soluble to a considerable extent in alcohol
and ether, forming brown solutions, and it is much more soluble in

FREE IODINE 305

carbon disulfide and in chloroform; all the iodine can be removed
from an aqueous solution by shaking a few times with either of these
solvents (cf. p. 17). Iodine is very soluble in hydriodic acid, or in a
solution of an alkali iodide, forming a triiodide:

, +I2.2 I

_ The solution of the triiodide shows all the reactions of free iodine;
but the vapor tension of the iodine is much less than in a solution
of pure water, because only a little free iodine is actually present at
any one time.

Commercial iodine always contains water, chlorine, bromine, and
often cyanogen (iodine cyanide) as impurities.

An aqueous soluticn of iodine is a weak oxidizing agent.

If hydrogen sulfide is passed through an aqueous solution of iodine,
it becomes colorless:and turbid, owing to the separation of sulfur:

H28+1,=2HI-+S.

Solid iodine is not acted upon at ordinary temperatures by hydrogen
sulfide; heat is necessary to produce the endothermic hydriodie acid.

In aqueous solution the necessary heat energy is furnished by the
solution of the hydriodic acid formed in water. The fact that solid
iodine is not acted upon by hydrogen sulfide, while it decomposes
arsine, is utilized in the preparation of hydrogen sulfide containing no |
arsenic from pyrites containing arsenic (cf. p. 234). The mixture
of hydrogen sulfide and arseniuretted hydrogen is passed over iodine,
and the latter only is removed.

Sodium Thiosulfate decolorizes iodine solutions, forming sodium
tetrathionate and sodium iodide:

282037 +Ise — S406 +21-.

Chlorine and bromine react in exactly the same way upon sodium thio-
sulfate when they are not present in excess. In the latter case the reaction
goes further and the tetrathionate is oxidized to sulfate and sulfuric acid with
deposition of sulfur, and the sulfur itself is gradually oxidized to sulfuric acid
by the halogens:

28,0¢-"+12H.0+8Cl, — 16Cl-+680,-+24Ht+S,,
S.+8H,.0+6Cl, =? 16Ht+280,-+12CI-.

Other weak oxidizing agents, such as ferric and cupric salts, act upon
thiosulfate similar to iodine (see Thiosulfuric acid).

Starch Paste.—Free iodine colors starch paste blue, but only in
the presence of hydriodic acid or a soluble iodide.

306 REACTIONS OF THE ACID CONSTITUENTS

Opinions differ concerning the composition of the blue “ iodide of starch.”
Some hold that it is a compound, while others regard it as a solid solution.*
According to Mylius,{ iodide of starch is the hydriodic acid compound of an
iodine addition-product of starch, containing about 18 per cent iodine, corre-
sponding to the formula [CaHO2l],-HI. This compound acts as an acid.
If iodide of starch is produced in a neutral solution, in the presence of iodides,
a salt of the above acid is formed, of which Mylius isolated the barium salt.
Iodide of starch, then, can be regarded as a double salt, similar to carnallite,
MgCl,-KCl1-6H,O. In dilute solutions it must be dissociated into its com-
ponents, e.g., the potassium salt

[CoH aoOvol] -KI = 4[CosH4oO201]+ KI,

and if we assume that the compound CH, OzI is colorless, the aqueous
solution of starch iodide will be colorless; but on increasing the concentration
of the alkali iodide, the double salt will be less dissociated and the blue color
of the undissociated compound will appear, which corresponds with the facts.
If a dilute aqueous iodine solution (obtained by shaking iodine with water)
is added drop by drop to a dilute aqueous starch solution, a blue color will
appear at the spot where the two solutions first come in contact, but this color
will disappear on stirring. If some potassium (or other) iodide is added to the —
colorless solution of the starch and iodine, a permanent blue coloration will
at once appear.{

The temporary appearance of the blue color, immediately on adding the
iodine solution, is probably due to the fact that the iodine first forms a sub-
stitution product with the starch, setting free hydriodic acid, which furnishes
the conditions for the formation of the iodide of starch.

The fact that a starch solution containing an iodide is much more sensitive
than one in pure water has been known for a long time,

Detection of Hydrochloric, Hydrobromic, and Hydriodic Acids in the
Presence of One Another 3

METHOD A

The solution to be tested should contain the alkali salts of the
above acids. Half of this solution is taken for the bromide and iodide
test, while the other half is retained for the chloride test.

(a) Detection of Bromine and Iodine

1. Acidify the solution with dilute sulfuric acid, add a little carbon disul-
fide, or chloroform, a drop of chlorine water and shake. If iodine is present

*Kisrer, Ann. 288, 689 (1894), C. O. Harz, Chem. Zentr., 1898, I, 1018;
AnpreEws and Gérscu, J. Am. Chem. Soc., 24, 865 (1906); Papra and Savort; Chem.
Zentrabl., 1905, I, 1593; Karyama, Z. anorg. Chem., 56, 209 (1907).

t Myruius, Ber., 20, 688, and C. Lonnus, Z. anal. Chem., 38, 409.

t The blue color disappears on heating the solution, but reappears on cooling.

FREE IODINE 307

ae in the presence of bromine), the carbon disulfide will be colored reddish-
violet.

To detect bromine shake repeatedly after the addition of small portions
of chlorine water, until the reddish-violet color has disappeared, showing that
the iodine has been completely oxidized to iodie acid; the brown color of the
bromine dissolved in the carbon disulfide will then appear and become a pale
yellow on further addition of chlorine water.

2. Instead of using chlorine water, it is often better to test for iodine
(especially when only small amounts are present, as in mineral waters) with
nitrous acid. Slightly acidify the solution to be tested for iodine and bromine
with dilute sulfuric acid, add carbon disulfide and a few drops of a solution
of nitrous acid in sulfuric acid and shake the mixture. If the carbon disulfide
is colored reddish-violet, iodine is present. Pour off aqueous solution (through
a moistened filter, in order to remove any suspended drops of colored carbon
disulfide), add chlorine water and shake the solution with fresh carbon disul-
fide. If the latter now assumes a brown color, bromine is present.

(b) Detection of Chlorine

The simplest way of separating chlorine from bromine and iodine is by
fractional precipitation with silver nitrate. If the solution containing salts
of the three halogens is treated with dilute silver nitrate, drop by drop, the iodine
will be first precipitated as yellow silver iodide, then the bromide as a pale
yellow silver salt, and finally the chlorine as pure white silver chloride. To
separate the chlorine from the other two halogens, acidify a little of the solution
to be tested with nitric acid, add a drop of dilute silver nitrate solution (1 : 100)
shake and boil the mixture, which causes the precipitate to collect together.
If bromine or iodine is present, the precipitate is yellow. Filter off the pre-
cipitate and again treat with the dilute silver nitrate, etc., until a pure
white precipitate of silver chloride is obtained, in case chlorine is present.

By filtering off the precipitate it is easy to tell whether it is pure white,
for the slightest tinge of yellow will show against the white paper.

METHOD B

This method of analysis is based upon the different degrees of readiness
with which the iodide, bromide and chloride ions are oxidized by 1 per cent
potassium permanganate solution. In a dilute solution containing a little
acetic acid and considerable sodium acetate, an iodide is immediately oxidized
with liberation of iodine, which can be removed from the solution by shaking
with a solvent which is immiscible with water (cf. p. 17). Under these
conditions an appreciable quantity of bromine or chlorine ions is not oxidized
in the time required for the necessary operations with an iodide. When the
concentration of the hydrogen ions is increased, by the addition of a prescribed
‘quantity of sulfuric acid, the bromide is oxidized very rapidly while the rate
of the corresponding reaction with chlorine ions is so slow that scarcely any
free chlorine is formed. The solution may even be boiled without losing more
than a small fraction of the chlorine ions originally present.

The increase in the oxidizing power of permanganate upon the addition of
hydrogen ions is a direct fulfillment of the mass-action principle (p. 13). In

308 ~ REACTIONS OF THE ACID CONSTITUENTS

the reaction between permanganate and halogen ions, hydrogen ions are
required, :
2Mn0,-+61-+8H* — 2Mn0,+4H30+38h,
or
2Mn0,-+101-+16H* — 2Mnt ++8H.0+5l:.

The permanganate will be reduced only to the quadrivalent condition, and
manganese dioxide will be precipitated, when the supply of hydrogen ions
is limited, but bivalent manganous cations will be formed to a greater extent
when more acid is present.

In carrying out this method of analysis, it is very important that the direc-
tions should be closely followed as regards the quantities of acid added.

Procedure.—If the solution to be tested is acid, add sodium carbonate
solution drop by drop until the solution no longer gives a decided red color to
blue litmus paper. If too much sodium carbonate is added accidentally, add
dilute nitric acid, drop by drop, until a very faint acid test is obtained in the
well-stirred liquid. Then add 8 ce. of normal sodium acetate solution, 2 ce.
of 6-normal acetic acid, and filter if any precipitate is formed. Add 3 ce. of

chloroform and 1 per cent permanganate solution, 1 cc. at a time, until the ~

aqueous solution after being vigorously shaken shows the pink color of per-
manganate. If an iodide is present, the chloroform will not be colored purple
by the presence of free iodine.

Pour the mixture through a moistened filter to remove the chloroform and
precipitated manganese dioxide; the wetting with water serves to prevent
chloroform passing through the pores of the paper. Shake the filtrate, if
necessary, with fresh portions of chloroform to extract all of the iodine from
the aqueous solution (cf. p. 17), filtering through a fresh filter each time.

When all the iodine has been removed, transfer the solution to a separatory
funnel, add 3 cc. of fresh chloroform, 5 cc. of 6-normal sulfuric acid and 1 ce.
of the dilute permanganate, unless an excess of the last mentioned reagent
is already present. Shake vigorously and then allow the chloroform to separate
out beneath the water. If bromine is present, the chloroform will be colored
yellow or orange. ,

Carefully remove the chloroform with the aid of the separatory funnel and
- transfer the aqueous solution to a porcelain dish. Add from 5 to 20 ce. of dilute
permanganate, according to the amount of bromide probably present, and boil
the mixture about five minutes, or until the volume of the solution has been
reduced to 10. cc. Filter off the manganese dioxide precipitate; and, if the
solution is still pink, add sulfurous acid drop by drop until it is colorless. Boil-

ing the solution serves to oxidize the last traces of bromine ions present with-

out oxidizing chlorine ions to any extent. The excess of permanganate reacts
with manganous ions (formed by reduction of the permanganate) to form
manganese dioxide,

2Mn0,-+3Mnt ++9H.0 — 5MnO:+4Ht,

but the presence of hydrogen ions interferes with this reaction, in accordance ~

with the mass-action law. Consequently, either owing to the fact that there
is a deficiency of manganous ions in the solution or because the concentration of

the hydrogen ions is too great, it is often necessary to add sulfurous acid to
reduce the last traces of permanganate ion,

2Mn0,-+580;-+6H*t — 2Mnt+-+580,-+-3H.0.

FREE IODINE 309

Dilute the solution to 100 cc., filter if necessary, add 20 ec. of 6-normal
nitric acid and a little silver nitrate solution. A curdy precipitate of silver
chloride is formed if a chloride is present.

Remark.—The above method of separation is capable of giving excellent
results. A solid substance can be tested for chloride, bromide and iodide after
first heating it with phosphoric acid and collecting the distillate. If about
2 gms. of solid, 25 cc. of water and 10 cc. of 85 per cent phosphoric acid are used,
all the iodine, bromine and chlorine will pass into the distillate. If any oxidizing
agent is present, however, free iodine, free bromine and free chlorine will be
formed in the order named. The free halogen can be removed by shaking the
distillate with chloroform. Iodine and bromine are recognized by the color
they impart to the chloroform (cf. p. 307). If free chlorine is formed, it will
dissolve in the chloroform and liberate iodine from potassium iodide. After
the removal of the free halogen, the distillate may be tested for hydriodie,
hydrobromic and hydrochloric acids as described above. By distilling with
phosphoric acid, halogen is obtained even from chlorates and bromates if a
reducing agent is present.

Detection of Halogen in the Presence of Cyanide

Conduct carbon dioxide through the slightly alkaline solution until the
escaping gas no longer gives any turbidity when passed into a nitric acid solu-
tion of silver nitrate.

The carbon dioxide may be prepared in a Kipp generator from marble and
dilute hydrochloric acid and washed with sodium bicarbonate solution. The
carbonic acid expels the weaker hydrocyanic acid from its salts (ef. p. 10).
Owing to the very poisonous nature of hydrocyanic acid, the expulsion of the
gas should take place under a good hood.

After the removal of the hydrocyanic acid, test the solution in the usual

way for halogens,

HYDROCYANIC AcID (PRussic AciD), HCN

Occurrence.—The compound of hydrogen with the univalent radi-
eal cyanogen, CN, occurs to some extent in nature. I¢ is found in
all parts of a tree growing in Java (Pangium Edule), particularly in
the seed kernels. Its compounds are found in many plants as a
glucoside (amygdalin), which yields, on hydrolysis, a carbohydrate,
benzaldehyde, and prussic acid:

CooH27NOi11+2H20 = 2C6H1206+CeHsCHO+HCN.
Amygdalin Benzaldehyde

This amygdalin is found in bitter almonds, in the fruit kernels
of cherries, apricots, peaches, etc., and in the leaves of the common
laurel tree (Prunus Laurocerasus).

Amygdalin is usually accompanied by a ferment, so that, on macer-
ating the parts of the plant which contain the amygdalin, an aqueous
solution of prussic acid is obtained (bitter-almond water).

Preparation.—If ammonia is passed over glowing carbon, ammo-
nium cyanide is formed; so that this salt, as well as other cyanogen
compounds, is found in the ‘“ gas-water’’ obtained by the dry dis- —
tillation of coal. ,

Hydrocyanic acid may also be prepared by adding acid to many.
cyanogen compounds. If yellow prussiate of potash is treated with
dilute sulfuric acid and distilled, prussic acid is evolved, which, after
standing over solid calcium chloride, may be obtained in the anhy-
drous condition as a colorless, exceedingly poisonous liquid, smelling
of bitter almonds, and boiling at 26.5° C.:

2K4[Fe(CN)«]+6H2S04 > 6KHSO14+KeFe[Fe(CN)¢]+-GHCN.

The best method for preparing anhydrous hydrogen cyanide *
consists in allowing a mixture of equal volumes sulfuric acid and
water to drop upon sticks of 98 per cent potassium cyanide. Hydro-
gen cyanide is evolved, contaminated with traces of water which can —
be removed by allowing the liquid to stand over solid calcium chloride.

Properties.—The liquid (as well as the gaseous) hydrocyanic acid

* J. Wave and L. Pantina, Proc. Chem. Soc., 190, 49 (1897-8).
310 F

HYDROCYANIC ACID — 311

burns with a reddish flame, and mixes in all proportions with water,
alcohol, and ether.

Aqueous hydrocyanic acid cannot be kept indefinitely; a brown
deposit soon appears, and ammonium formate is formed:

HCN+2HOH = HCO2NHsg.

If a little mineral acid is added to the aqueous solution, it may
be kept much longer; but, even then, ammonia and formic acid will
be formed after a long time. Cold concentrated hydrochloric acid
converts hydrocyanic acid into formamide, HCN+H20=HCONH»p,
but on warming this compound is decomposed into carbon monoxide
and ammonia. :

Hydrocyanic acid in aqueous solution is a very poor conductor
of electricity; in other words, it is a very weak acid, and is ionized
only to a slight extent.

The salts of hydrocyanic acid, the cyanides, are very similar in
their properties to the corresponding halogen compounds, but are
distinguished by their ability to form stable complex salts, which
contain scarcely any cyanogen ions in aqueous solution, and conse-
quently do not give some of the reactions for hydrocyanic acid.

Solubility of Cyanides.—The cyanides of the alkalies and alkaline
earths are readily soluble in water, but hydrolysis (p. 48) takes place
to a considerable extent:

CN-+H20 — OH-+HCN.

Since hydrocyanic acid is only slightly ionized, the aqueous solu-
tion of an alkali cyanide behaves like a solution of alkali hydroxide
containing free prussic acid; the smell of the latter can be readily
detected.

The remaining cyanides, with the exception of mercuric cyanide,
are insoluble in water.

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid decomposes solutions of all soluble cyan-
ides, with the exception of mercuric cyanide, setting free hydrocyanic
acid in the cold (recognizable by its odor). Danger!

The insoluble cyanides are decomposed by dilute sulfuric acid
only on warming.

2. Concentrated Sulfuric Acid decomposes all cyanides on warm-
ing, the complex cyanides as well as the simple ones. The metals
are then obtained in the form of acid sulfates, the carbon of the

312 REACTIONS OF THE ACID CONSTITUENTS

eyanogen is changed to carbon monoxide, and the nitrogen into am- .
monium sulfate:
> Ni(CN)2+2H2S04+2H20 = Ni8SO4+ (NHa)2804+2CO ts

With mercuric cyanide, besides carbon monoxide, sulfur dioxide
and carbon dioxide are obtained; for mercuric cyanide is decomposed
at the temperature of boiling sulfuric acid into mercury and cyanogen;
and the former dissolves in the hot sulfuric acid, with formation of
mercuric sulfate and evolution of sulfur dioxide:

Hg(CN)2+6H2S804 — 2NH4HS01+ Hg( A804 '2
+CO Tt +COz2 T +802 | +803.

3. Silver Nitrate.—If silver nitrate is added to a solution of an
alkali cyanide drop by drop, a precipitate is formed on the addition
of each drop, which, however, redissolves on stirring the liquid, for
silver cyanide is soluble in an excess of alkali cyanide:

CN-+Agt > AgCN;
AgCN-+4ON- > [Ag(CN)oI-.

The complex silver cyanide ion is decomposed by further addition
of silver nitrate, being finally completely transformed into insoluble
silver cyanide:

[Ag(CN)2]--+Agt > 2AgCN.

Consequently the precipitation is complete only when an excess
of silver nitrate is added. |

Silver cyanide is insoluble in water and dilute nitric acid, percept-
ibly soluble in concentrated nitric acid, and readily soluble in ammonia,
sodium thiosulfate, and potassium cyanide. Dilute nitric acid repre-
cipitates silver cyanide from the solution in ammonia or potassium
cyanide.

Concentrated hydrochloric acid decomposes silver cyanide, on
warming, into silver chloride, with evolution of hydrocyanic acid
(difference from silver chloride, bromide, or iodide).

By igniting silver cyanide, there is formed cyanogen gas, metallic
silver, and brown, difficultly volatile paracyanogen, which, on further
heating, is completely volatilized, leaving behind pure silver: ;

2AgCN =2Ag+-(CN)o.

Much more certain than the silver nitrate test are the tests for
hydrocyanic acid which depend upon the formation of Prussian blue
or of ferric thiocyanate.

HYDROCYANIC ACID 313

4. Prussian Blue Reaction.—Prussian blue is formed by the action
of ferric salts upon potassium ferrocyanide (cf. p. 150).

3[Fe(CN).6]-~+4Fet tt > Fe4[Fe(CN) sls.

In order, therefore, to apply this reaction to potassium cyanide, etc., it is
necessary first to transform the cyanide into potassium ferrocyanide. This
may be accomplished by the addition of a ferrous salt, whereby ferrous eyanide
is first formed, which dissolves in an excess of potassium cyanide, forming
potassium ferrocyanide:

(a) Fe+++2CN- — Fe(CN),;
(b) Fe(CN):+4CN- > [Fe(CN),]->.

Potassium ferrocyanide is formed even more readily by the action of potas-
sium cyanide upon ferrous hydroxide:

Fe(OH).+2CN~ — Fe(CN).+20H_;
Fe(CN),:+4CN- — [Fe(CN)«]=~.

For the formation of the potassium ferrocyanide a little iron and considerable
potassium cyanide are required. Consequently, to bring about the reaction, add
a little ferrous sulfate to the alkaline solution of an alkali cyanide and boil the
mixture. Then add a little hydrochloric acid, whereby a clear solution
is obtained, which gives, with a little ferric chloride, the blue precipitate. If
only traces of hydrocyanic acid are present, the solution appears green at first,
but, after standing some time, “ flocks ” of Prussian blue will be precipitated.

The Prussian blue reaction is used for the detection of nitrogen in organic
substances (cf. p. 316).

Remark. —When only a trace of cyanide is present, the Prussian blue seen
through yellow ferric chloride appears green. If the ferric chloride solution is
diluted largely to get rid of the yellow color, the iron is then largely in colloidal
solution, as a result of the formation of basic salt by hydrolysis; such a solution
reacts very slowly with ferrocyanide. A more sensitive test for traces of
ferrocyanide is obtained by adding a saturated solution of ferrous sulfate.
Such a solution always contains enough ferric ions to give the Prussian blue test
with a trace of ferrocyanide ions. When more than a trace of ferrocyanide is
present, however, it is best to test with ferric chloride solution as outlined above.

5. The Ferric Thiocyanate Reaction.—Potassium thiocyanate pro-
duces a red coloration with a ferric salt, owing to the formation of
slightly ionized ferric thiocyanate (cf. p. 150):

3CNS-+Fet + *— Fe(CNS)s,

The cyanide, therefore, must be changed to thiocyanate in
order to apply this reaction, which can be done by heating with sulfur,

KCN+S=KCNS,

314 REACTIONS OF THE ACID CONSTITUENTS

or, better, by treatment with an alkaline polysulfide,
KCN+(NHa4)2S2 = (NHa)25+KCNS.

To the concentrated solution of the cyanide (in a porcelain dish) add a
little yellow ammonium sulfide and evaporate the mixture to dryness
on the water-bath. Add a little hydrochloric acid and a drop of ferric chloride
solution; the characteristic blood-red coloration will be produced if only traces
of cyanide are present. .

It is necessary to acidify in order to destroy the (NH4)25, which would
precipitate black Fe.S; with FeCl; and the red coloration would not appear.

|
6. Mercurous Nitrate produces a gray precipitate of metallic
mercury when added to a solution of an alkali cyanide (difference

from a chloride, bromide, or iodide):
Hg.t ++2CN~ — Hg(CN)2+Hg.

Behavior of Mercuric Cyanide

Mercurie cyanide, Hg(CN)s, is a non-electrolyte, is quite soluble in water,
in alcohol and in ether, and behaves quite differently from all other, cyanides,
All the above-mentioned reactions, with the exception of the ferric thiocyanate
test, fail with this compound. Mercuric cyanide solution gives no precipitate
with silver nitrate, but a readily soluble double salt is formed,

Similarly there is no precipitate formed with ammonia, potassium iodide, alkali.
hydroxide, or alkali carbonate because all these reagents give, under ordinary
conditions, precipitates of mercuric compounds which are soluble in potassium
cyanide. Thus, for example, mercuric oxide dissolves easily in potassium
cyanide:

HgO+2KCN+H,0 =2KOH+He(CN)s.

Mercuric oxide itself is fairly soluble in mercuric cyanide:

_o/ He—CN
HgO+Hg(CN)2 = OX He CN:

Hydrochloric, hydrobromic and hydriodic acid decompose mercuric cyanide
readily, but dilute sulfuric acid alone has little action upon it. In the presence
of a soluble chloride, however, mercuric cyanide is decomposed easily by sul-
furic acid, or even by oxalic acid or tartaric acid. If, therefore, a solution of
mercuric cyanide is treated with common salt and dilute sulfuric, oxalic or tar-
taric acid, then, on subjecting the mixture to distillation, a distillate is obtained
containing hydrocyanic acid, showing all the tests characteristic of this acid.

Mercuric cyanide in solution is acted upon by hydrogen sulfide or by alkali
sulfides, forming mercuric sulfide and hydrocyanic acid, or one of its salts. If
the black precipitate is removed by filtration, the ferric thiocyanate test is
obtained without difficulty.

HYDROCYANIC ACID 315

Behavior of Cyanides on Ignition

The cyanides of the alkalies and alkaline earths fuse without decomposition
when heated out of contact with the air; heated in contact with air, they absorb
oxygen with avidity, forming cyanates:

2KCN +0, =2KCNO.

Consequently the alkali cyanides are strong reducing agents (cf. p. 260).

The cyanides of the bivalent heavy metals are decomposed on ignition, out
of contact with the air, into nitrogen and metallic carbide; the latter often being
further decomposed into metal and carbon:

Fe(CN), =FeC.+N, T ;
Pb(CN): =Pb+2C+N; Tf.

The cyanides of the trivalent metals are unknown in the free state; those

of the noble metals are decomposed, by igniting, into metal and dicyanogen:
Hg(CN):=Hg+(CN): T.
It is a characteristic property of the cyanides of the heavy metals that they

are readily soluble in alkali cyanide solutions, forming very stable complex com-

pounds, which are to be regarded as salts of the following acids: ,

H[R‘(CN).], H.[R™4(CN).], Hs[R™(CN).] and H.{R™(CN),].
The first two of the above acids are so unstable that they are decomposed,
as soon as they are set free, into hydrocyanic acid and cyanide:
» H[R(CN).] =HCN+RCN;
H.[R(CN).] =2HCN+R(CN)>..

Consequently all cyanides which are derived from these acids evolve hydro-
cyanic acid when treated with dilute hydrochloric or sulfuric acid in the cold.
Such compounds are:

K[AgCN], K[Au(CN)]., K2[Ni(CN),], Ke[Zn(CN),], Ke[Cd(CN),], ete.

These salts must be regarded as complex compounds (p. 24), for their
aqueous solutions contain almost no heavy metal ions; they are not precipitated
by caustic alkali, alkali carbonate, or ammonia. From this fact it follows that
the oxides of these metals are soluble in cyanides of the alkalies, forming the

following complex salts:
AgO+4CN~+H,0 — 20H~+2[Ag(CN),]-;
Ni0+4CN~+H20 — 20H~-+[Ni(CN).]-
ZnO+4CN~+H,20 — 20H +[Zn(CN),]-;
Cd0+4CN-+H,20 — 20H~-+[Cd(CN),}-,

316 REACTIONS OF THE ACID CONSTITUENTS

Hydrogen sulfide decomposes the silver and cadmium salts without difficulty
the zine salt slightly, and the nickel salt not at all.

The acids of the general formula H,{R™(CN).] and H.[R™(CN),] are, in
contrast with the above acids, quite stable in the free state, and can be usually
obtained, without the loss of hydrocyanic acid, by acidifying a solution of one
of their salts with cold dilute mineral acid; but, on warming the acid solution,
hydrocyanic acid is given off.

As typical representatives of these acids we have ferrocyanic, ferricyanic
and cobalticyanic acids.* ‘

We will consider ferrocyanic and ferricyanic acids separately; but before
doing so we will briefly describe

DICYANOGEN

which is obtained by heating the cyanides of the noble metals, as a colorless
gas with a disagreeable odor; it burns with a reddish flame, and is soluble in
water (25 parts water dissolve 100 parts of dicyanogen). The aqueous solution
eannot be kept very long, as brownish “ flocks” separate out little by little
(azulmie acid, C,H;N;O) and the solution then contains ammonium cyanide,
ammonium carbonate, ammonium oxalate, and urea.

Just as chlorine acts upon caustic alkalies, forming chloride and hypochlorite,
so dicyanogen reacts with them, forming a cyanide and a cyanate:

Cl.+2KOH =KCI+H:.0+ KOCI;
(CN)2.+2KOH =KCN+H:,0+KOCN.

On conducting hydrogen sulfide into a solution of dicyanogen, a red,
crystalline precipitate of hydrorubianic acid is obtained:

Cf. p. 219. 2 aoe abeadd =(CSNH2)2.

Detection of Nitrogen in Organic Substances. (Lassaigne.)

If a little nitrogenous, organic substance is placed with some metallic sodium
in a small glass tube which is closed at one end, and the tube is heated till the
glass begins to soften, then the nitrogen and some of the carbon of the organic
substance will combine with sodium to form sodium cyanide. After heating —
for two minutes, plunge the hot end of the glass tube in a little water, whereby
the glass breaks and the contents of the tube are immediately wet by water.
Filter off the alkaline solution containing the sodium cyanide from the residual
carbon and glass splinters, add a little ferrous sulfate solution and boil. Finally
add a few drops of ferric chloride solution and enough hydrochloric acid to
neutralize the alkali. If nitrogen was present in the original substance, Prussian _
blue is now obtained.

Remark.—Certain nitrogenous substances are decomposed with evolution of
nitrogen before’ the temperature required for the cyanide formation is reached, f

* Cobaltocyanic acid is extremely unstable, like mangano- and manganicyanic
acids. Its salts evolve HCN when treated with cold, dilute, mineral acids.
ft Grae, Ber., 17, 1178 (1884),

DICYANOGEN 317

and thus the nitrogen escapes the Lassaigne reaction; in other cases the experi-
ment fails on account of the extreme volatility of the organic substance.*

According to E. A. Kehrer,{ the Lassaigne reaction gives positive results
in all cases if the sodium is first heated by itself and then brought in contact
with the vapors of the organic material. The experiment is carried out in a
small tube, drawn out at the. closed end, such as used for the arsenic test (cf.
Fig. 21, p. 235). The substance is placed in the restricted part of the tube
and then, in the wider part, just before the restriction is reached, a small piece
of sodium is introduced which has been rolled between the fingers and freed from
petroleum by touching it to filter paper. The sodium is heated until it glows
and then, by means of another, small flame, the substance to be tested is heated
so that it melts and the vapors just rise up to the glowing sodium but hardly
reach through it. By taking away the small flame, the vapors can be made to
condense and be driven toward the glowing sodium again. Otherwise, the test
is carried out exactly as described above.

For the detection of traces of nitrogen it is best not to add any ferric
chloride, but to add a saturated solution of ferrous sulfate, after adding the
acid (cf. p. 313.)

Detection of Hydrocyanic Acid in the Presence of Halogen Acid,
Ferrocyanic, Ferricyanic and Thiocyanic Acids

Hydrocyanic acid is by far the weakest of all the above acids (cf. p. 10)
and it alone is expelled from its salts by means of carbonic acid at the tempera-
ture of boiling water. \

Place the solution to be tested in a small Erlenmeyer flask, add 0.6 to 1 gm.
of NaHCO; and close the flask with a two-hole rubber stopper. Insert through
one of the holes in the stopper a piece of glass tubing that reaches nearly to the
bottom of the flask and serves for the introduction of the carbon dioxide gas;
through the other hole insert a piece of tubing that reaches only to the bottom of
the rubber stopper and serves for the escape of the gas. Pass carbon dioxide
gas (cf. p. 309), through the liquid in the flask, gradually heat to boiling and
conduct the escaping vapors into silver nitrate solution which has been acidified
with nitric acid.

If cyanide is present in the original solution, a white precipitate of silver
cyanide is formed in the silver nitrate solution within a short time. To con-
firm the test, decant off the solution from the precipitate and wash it a few times
by decantation with water. Cover the precipitate with a little yellow am-
monium sulfide, warm, and filter. Evaporate the filtrate to a small volume
and treat with a little hydrochloric acid and a few drops of ferric chloride solu-
tion. A red color, due to ferric thiocyanate, proves the presence of cyanide.

* Frist, Ber., 35, 1559 (1902).
t Ber., 35, 2523 (1902).

FERROCYANIC ACID, Ha[Fe(CN).¢]

Ferrocyanic acid is a white, solid substance, which is readily
soluble in water and in alcohol, the solution soon becoming blue on
exposure to air. The salts of this acid are much more stable than the
acid itself, being all prepared from the potassium salt, the so-called
yellow prussiate cf potash. This potassium salt, the most important
ferrocyanide of commerce, is obtained by the fusion of organic sub-
stances containing nitrogen and sulfur (blood, etc.) with potash and
metallic iron, and by lixiviating the product of the fusion with water.

In the melt, iron sulfide and potassium cyanide are found, which,
on treatment with water, are changed to potassium ferrocyanide
and potassium sulfide,

FeS+6KCN = Ka[Fe(CN)¢]+K28,

and, on evaporating the solution, the former salt separates out (with
three molecules of water of crystallization) in the form of large, yellow,
tetragonal octahedrons.

Recently this salt has been obtained as a by-product in the manu-
facture of illuminating-gas, Prussian blue and ammonium thiocyan-
ate being formed from the purification of the gas.

The following equations will give some idea of the formation of
potassium ferrocyanide in the gas-house:

1. Fe7(CN)is+6Ca(OH)2=4Fe(OH)3 +3Cazg|Fe(CN)6] ;
2. Caz{Fe(CN)6]+2KCl = K2Ca[Fe(CN)¢6]+CaCh;

Very difficultly soluble
3. Ke2CalFe (CN) 6] oh KeCO3 = CaCO; + K4lFe (CN) 6] ‘

Solubility of Ferrocyanides—The ferrocyanides of the alkalies
and alkaline earths are soluble in water; but the remaining salts
dissolve with difficulty (if at all) in water and in cold dilute acids.

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid.—The ferrocyanides are not decomposed

by cold sulfuric acid, but break up at the boiling vemnperenyes with
evolution of ivdniesnaiie acid:

2Ka|Fe(CN) 6] +6H2S04 — K2Fe[Fe(CN)¢]-+6KHSO4+-6HCN.
318

FERROCYANIC ACID 319

2. Concentrated Sulfuric Acid decomposes ferrocyanides com-

pletely, on warming, with evolution of carbon monoxide, which burns
with a blue flame:

K4l[Fe(CN).6]+11H2SO04+6H20 =
= FeS0O4+4KHS01+6NHsHSO4+6CO 7 .

SO, is also liberated by this reaction, as a part of the ferrous sulfate is
oxidized by the sulfuric acid to ferric sulfate:

2FeSO,+280; “<> Fe:(SO,);+S0, T °
3. Silver Nitrate produces a white precipitate of silver ferrocyanide
[Fe(CN)6~~]+4Agt — Ags[Fe(CN)cl,

insoluble in dilute nitric acid and ammonia, but soluble in potassium
cyanide solution. On treatment with concentrated nitric acid, it is
changed to orange silver ferricyanide, and is then soluble in ammonia.

4. Barium Chloride gives no precipitation.

5. Ferric Salts produce a precipitate of Prussian blue in neutral
or acid solutions (cf. p. 150).

6. Ferrous Salts yield a light blue precipitation, which changes
to a darker blue on exposure to the air (cf. p. 147).

7. Cupric and Uranyl Salts produce brown precipitates.

To detect ferrocyanic acid in an insoluble ferrocyanide, boil the latter
with caustic alkali solution; metallic hydroxide and a ferrocyanide will be
formed. Thus, Prussian blue yields insoluble ferric hydroxide and a soluble
ferrocyanide:

Fe,{Fe(CN)«];+120H- — 4Fe(OH);+3[Fe(CN),]-—.

Filter off the insoluble hydroxide, add dilute hydrochloric acid to the filtrate

and treat with ferric chloride, Prussian blue is again formed if a ferrocyanide is

resent.

Prussian blue is often used in wall-papers as a pigment. If it is desired to
detect the presence of this compound in a wall-paper, cut about 100 cm*. of
the paper into small pieces, boil them with caustic potash solution, filter, and
treat the filtrate according to the method just described. In a few hours a
distinct blue precipitate of Prussian blue will be visible in the bottom of the
test-tube, if it was originally present.

Some insoluble ferrocyanides do not yield the hydroxide of the metal on
treatment with caustic alkali. Thus the brown uranyl ferrocyanide yields
insoluble yellow potassium uranate and soluble potassium ferrocyanide (cf.

; 188);
: Ce zine ferrocyanide is completely soluble in caustic alkali, forming
an alkali zincate and soluble ferrocyanide:

Zn,[Fe’’(CN).|+8OH~ — [Fe(CN).J-~+2Zn0.-+4H,0.

320 REACTIONS OF THE ACID CONSTITUENTS

In order to separate the zine from the ferrocyanide, pass carbon dioxide gas
into the solution, boil, and filter off the insoluble zine carbonate. The filtrate
then contains potassium ferrocyanide, which can be detected as above.

8. Lead Salts precipitate white lead ferrocyanide insoluble in
dilute nitric acid.

9. Thorium Nitrate added to a slightly acid solution of an alkali

ferrocyanide produces a white precipitate, difficult to filter (differ-
ence from ferricyanic and thiocyanic acids).

Behavior of Ferrocyanides on Ignition

On being ignited, the ferrocyanides yield iron carbide, cyanide,

and nitrogen:
Ka[Fe(CN)¢] = AKCN +FeCo+No 13

AgalFe(CN)¢] =4AgCN-+FeCo+Np 7.

In the latter case, the silver cyanide is further decomposed into
metal and dicyanogen:

2AgCN =Age+(CN)s f.

FERRICYANIC ACID, H3{Fe(CN)o]

Ferricyanic acid forms brown needles, readily soluble in water.

Tts salts, the ferricyanides, are very stable, and are obtained by
the oxidation of the corresponding ferrocyanides. The most important
of all these salts, potassium ferricyanide (red prussiate of potash),
K3[Fe(CN)6], is obtained by the oxidation of potassium mpi:
with chlorine:

2K4[Fe(CN) 6] +Cle =2KC1+2K3[Fe(CN)¢].

Bromine, hydrogen peroxide, etc., may be used instead of chlorine.

‘Solubility of Ferricyanides—The ferricyanides of the alkalies
and alkaline earths, and the ferric salt of ferricyanic acid, are soluble
in water, but the remaining salts are insoluble even in dilute acids.

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid evolves no hydrocyanie acid in the cold
(difference from cyanides), but does so on warming with the acid.

2. Concentrated Sulfuric Acid decomposes all ferricyanides, on
warming, with the formation of sulfates and carbon monoxide:

Ks3[Fe(CN)¢]+11H2S04+6H20 =
-=FeH(SOs)2+3KHS01+6NH4HS01+6C0 1 .
3. Silver Nitrate produces orange silver ferricyanide:
[Fe(CN)6]=+3Ag* > Ags[Fe(CN)g],

soluble in ammonia, but insoluble in nitric acid.
4. Barium Chloride gives no precipitation.
5. Ferrous Salts produce, in neutral and acid solutions, a pre-
cipitate of Turnbull’s blue (cf. p. 147).
6. Ferric Salts produce no precipitation, but a brown coloration.
7. Cupric Salts yield green cupric ferricyanide:

2 [Fe(CN)6]=+3Cut* — Cus[Fe(CN)elo.

8. Behavior of Ferricyanides in Alkaline Solutions.—Ferricyanic
acid is a strong oxidizing agent in alkaline solutions, being readily
reduced to ferrocyanic acid by sulfide, iodide, sulfite, ferrous

321

322 REACTIONS OF THE ACID CONSTITUENTS

hydroxide, manganous hydroxide, lead oxide, starch, cellulose (paper),
etc.; e.g.:

2[Fe(CN)6]" +S" — 2[Fe(CN)e~" +8;
2[Fe(CN)o]"+2I- — 2[Fe(CN)6]~ +e;
2{Fe(CN).6]=+S03"-+20H- — 2[Fe(CN)6]/>~ +S04-+H20;
2{Fe(CN)¢6]=+PbO2- — 2[Fe(CN)6]” +PbO2;
[Fe(CN)o]"+Fe(OH)2+OH™ — [Fe(CN)6]~~+Fe(OH)s.
The ferricyanides are even reduced by ammonia, forming nitrogen:
6[Fe(CN)6]=+8NH3 — 6[Fe(CN)¢l7~+Ne T +6NHa4".

On account of this easy reducibility of ferricyanic acid, it is
often difficult, sometimes impossible, to detect its presence, par-
ticularly in an insoluble compound. If Turnbull’s blue is boiled
with caustic potash, the residue will consist, of a mixture of ferrous
and ferric hydroxides, and the solution will contain potassium fer-
rocyanide. See p. 147.

The behavior of cyanides toward suspended, yellow mercuric oxide

is very important. Almost all cyanides, simple or complex, with the.
exception of potassium cobalticyanide, are completely decomposed —

by this reagent. Mercuric cyanide and an oxide of the other metal
are formed, which, if insoluble, may be separated from the mercuric
cyanide by filtration. Thus potassium ferrocyanide is decomposed
by mercuric oxide as follows:

K4[Fe(CN) 6] +3Hg0O +3H20 = Fe(OH)2+4KOH+3Hg(CN)z.
Prussian blue as follows: |
Fe4[Fe(CN)¢!3 +9HgO+9H20 = 8Fe(OH)2+4Fe(OH)3 +9Hg(CN)e.
This decomposition of the cyanides by mercuric oxide is often

used in quantitative analysis for the separation of metallic cyanides.

Behavior of the Ferricyanides on Ignition

The ferricyanides are decomposed into iron carbide, cyanide,

dicyanogen, and nitrogen:
2K3[Fe’"(CN)¢] =2FeC2+6KCN+2N2 T +(CN)e T.

By heating a ferricyanide in a closed tube, dicyanogen therefore
is given off, which burns with a reddish flame.

THIOCYANIC AcID, HCNS

Thiocyanic acid is found in small amounts, in the form of its sodium
salt, in saliva and urine.

The free acid is a colorless, unstable liquid, with a penetrating
odor. It can be kept better in aqueous solution than in the anhydrous
state, but its salts, the thiocyanates, are much more stable than the
acid itself. The alkali salts can be prepared from the corresponding
cyanides by heating with sulfur:

KCN+S=KCNS.

They may also be prepared by treating hydrocyanic acid or an
alkali cyanide with an alkali polysulfide at ordinary temperatures,

KCN -+(NH4)2S2 = (NH4)28+KCNS,

or by boiling an alkali thiosulfate solution with an alkali cyanide
(cf. p. 389):
Na2S203-+KCN = NasSO3+KCNS.

The easiest way to prepare ammonium thiocyanate is to allow
a mixture of 30 cc. concentrated ammonium hydroxide, 30 ce. alcohol:
and 7 ec. carbon disulfide, to evaporate very slowly on the water bath.
First of all, ammonium thiocarbamate is formed

CSe+2NH3 = NH4CS2N Hae,
and during the evaporation this loses hydrogen sulfide:
NH.iCS2NH2=NH4CNS+H2S.

Solubility—Most of the thiocyanates are soluble in water; excep-
tions are the silver, mercury, copper, and gold salts. Lead thiocyanate
is difficulty soluble in water; on boiling with water it is decomposed.

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid (double normal) causes no reaction.

2. Moderately Concentrated Sulfuric Acid (14-normal) decom-
poses the thiocyanates, with evolution of carbonyl sulfide, which burns
with a blue flame:

KCNS+2H2S04+H20 = KHSO4+ (NH4)HSO4+CoOS 7 .
323

324 REACTIONS OF THE ACID CONSTITUENTS

3. Concentrated Sulfuric Acid violently decomposes thiocyanates,
with evolution of very disagreeably smelling vapors, COS, HCOOH,

COz, SOe, and deposition of sulfur.
4. Silver Nitrate precipitates white, curdy, silver thiocyanate,

CNS~+Ag* — AgCNS,
insoluble in dilute nitric acid, soluble in ammonia.

5. Ferric Salts produce a blood-red coloration, due to the for-
mation of non-ionized ferric thiocyanate,

3CNS”-+Fe — Fe(CNS8)s,

very soluble in ether (cf. p. 150).
6. Mercuric Nitrate precipitates white mercuric thiocyanate,

Hg*+++2CNS8-+Hg(CNS)>,

very difficultly soluble in water, but readily soluble in an excess of
potassium thiocyanate:

Hg(CNS)2+CNS” — [Hg(CNS)s3]-.

If dry K[Hg(CNS)s] is heated, the salt expands greatly (Pharaoh’s .
serpents).
7. Mercuric Chloride gives a precipitate only after long standing. —
8. Mercurous Nitrate produces a gray to black precipitate. On
adding mercurous nitrate drop by drop to a fairly concentrated solu-
tion of potassium thiocyanate, a gray precipitate of metallic mercury
is first obtained, and the solution contains potassium mercuric thio-

cyanate:
Hgo**++3CNS~ — [Hg(CNS)3]-+Hg.

If the addition of mercurous nitrate is continued until no more mer-
cury is precipitated, and the solution then filtered, the filtrate will
contain potassium mercuric thiocyanate; but, on adding still more
mercurous nitrate, pure white, mercurous thiocyanate is precipitated:

2{Hg(CNS)s3]-+3Hg2* * — 2Hg*++3Hg2(CNS)s.

If, on the other hand, a very dilute solution of potassium thio-
cyanate is added to a very dilute solution of mercurous nitrate, the
white precipitate of mercurous thiocyanate is obtained directly: _

Hg2t*+2CNS~ — Hg2(CNS)p.

9. Cupric Salts.—On adding a few drops of a solution containing
a cupric salt to one of an alkali thiocyanate, the solution is colored

THIOCYANIC ACID 325

emerald-green; and, on further addition of the copper solution, black
cupric thiocyanate is precipitated. If sulfurous acid is added,
white cuprous thiocyanate is deposited,

2Cut++S03"+2CNS-+H:0 — Cu2(CNS)2+2H*+80¢,

insoluble in dilute hydrochloric and sulfuric acids.

10. Cobalt Salts.—If a solution containing an alkali thiocyanate
is treated with a small amount of a cobalt salt, and the solution
shaken with a mixture of equal parts amyl alcohol and ether,* the
upper layer of alcohol ether separates out azure-blue in color (ef.
p. 182). This reaction is. analogous to that of cyanic acid upon cobalt
salts (cf. p. 344).

Detection of Thiocyanates in the Presence of Halogen and Cyanide

| First free the solution from hydrocyaniec acid by adding a little sodium
bicarbonate, heating to boiling and passing carbon dioxide gas through the
solution until the escaping gas led into slightly acid silver nitrate solution
gives no turbidity of silver cyanide.

After the removal of the hydrocyanic acid, test for thiocyanate, in the
absence of iodide, by acidifying with hydrochloric acid and adding a few drops
of ferric chloride. A blood-red coloration shows the presence of thiocyanate. The
test cannot be obtained satisfactorily in the presence of an iodide because it
also will react with ferric chloride, causing liberation of free iodine.

If an iodide is present, add a little nitric acid to the solution from which
the hydrocyanic acid has been expelled, and precipitate the halogens and
thiocyanate by the addition of an excess of silver nitrate solution. After the
precipitate has settled, decant off the supernatant solution and wash the
precipitate several times by decantation with water. Then shake the precipi-
tate vigorously with 6-normal ammonia solution. This dissolves the chloride
and thiocyanate readily, all or a part of the bromide, but no appreciable quantity
of silver iodide. Filter and add colorless ammonium sulfide to the filtrate.
Filter off the silver sulfide precipitate, add a drop of sodium carbonate solution,
evaporate to small volume, acidify with hydrochloric acid, and test for thio-
cyanate with ferric chloride solution.

Detection of Halogens in the Presence of Thiocyanate
VoLHARD’s MrEetTHop

Treat the nitric acid solution with an excess of silver solution, filter and
dry the precipitate with suction. Transfer the precipitate to a porcelain dish
and heat it on the water bath for 45 minutes with concentrated nitric acid.
This causes the complete decomposition of all the thiocyanate,

GAgCNS+4H,0+16HNO; =3Ag,S0,+3(NH,).S0.,+5CO, T +16NO 1,

* Or with amyl alcohol alone.

’

326 REACTIONS OF THE ACID CONSTITUENTS

and does not attack the silver halides appreciably. Dilute, filter and wash a
few times with hot water. Reduce the precipitate with zinc and dilute sulfuric
acid (p. 281) and test the filtrate according to p. 307.

Or, instead of treating the washed silver precipitate with nitric acid it may
be boiled with sulfuric acid (1 : 1) until the precipitate becomes black and col-
lects in a ball:

2AgCNS+2H.80,+3H.0 =2NH,HSO.+COS Tf +CO: T +AgS.

Dilute the solution with water, filter off the silver halides and silver sulfide,
wash with water, reduce with zinc and sulfuric acid and test for the halogens —
after boiling off the hydrogen sulfide from the last filtrate.

Remark.—It is always necessary to destroy the thiocyanate before reducing
with zinc and sulfuric acid, because otherwise hydrocyanic acid will be formed:

2AeSCN-+3H, > 2Ag-+2HCN-+2H,S 1.

According to A. W. Hofmann, free thiocyanic acid on being reduced with
nascent hydrogen yields a mixture of thioformaldehyde, methylamine, ammonia
and hydrogen sulfide.

Testing Commercial Alkali Thiocyanate for Chloride
C. Mann’s MetHop

Dissolve 5 gms. of the alkali thiocyanate in 20 cc. of water and to the solu-
tion add 20 gms. of crystallized copper sulfate dissolved in 100 cc. of water.
A black precipitate of cupric thiocyanate is formed:

Cutt++2CNS- — Cu(CNS),

Pass hydrogen sulfide gas through the solution until the precipitate becomes
nearly white:

2Cu(CNS),-+H.$ =Cu(CNS).+8+2HCNS.

‘Then, when the supernatant blue copper solution begins to get brown,
owing to the formation of copper sulfide, stop introducing the hydrogen sulfide
gas and allow the liquid to stand a few hours. During this time the thiocyanic

acid formed by the above reaction is acted upon by the copper sulfide present,
as follows:

2Cu(CNS).+2CuS — 2Cu,(CNS).+28.

Filter and treat the filtrate with silver nitrate. A white precipitate shows
chloride to be present.

Remark.—An equally good method is the following: Treat the solution of
the thiocyanate with an excess of copper sulfate solution and-introduce sulfur
dioxide gas until the precipitate becomes white cuprous thiocyanate. Allow
the solution to stand several hours, then filter off the cuprous thiocyanate, treat
the filtrate with nitric acid and test for chlorine with silver nitrate.

Instead of sulfurous acid, hydroxylamine sulfate may be used to reduce the
cupric solution.

THIOCYANIC ACID 327

Detection of Ferro- and Ferricyanides in the Presence of Thiocyanate
Mertuop or P. E. Brownine anp H. E. Parmer

Acidify the dilute solution of the alkali salts of these acids with acetic —
acid or hydrochloric acid, avoiding an excess, and add a solution of thorium
nitrate; finely divided thorium ferrocyanide will be precipitated. Shake the
solution with finely divided asbestos, or filter paper pulp, filter and wash the
precipitate with cold water. Pour dilute sodium hydroxide solution over the
precipitate, acidify the solution thus obtained and add a few drops of ferric
chloride; Prussian blue is formed if a ferrocyanide was present.

Add cadmium sulfate solution to the filtrate from the thorium ferrocyanide
precipitate, shake with finely divided asbestos and filter off the cadmium ferri-
cyanide. Wash the precipitate with cold water, dissolve in sodium hydroxide
solution, acidify the solution with hydrochloric acid and treat with ferrous sul-
fate solution. The formation of Turnbull’s blue shows the presence of ferricyanic
acid.

Adding ferric chloride to the filtrate from the cadmium ferricyanide pre-
cipitate, a blood-red coloration will be obtained if a thiocyanate is present.

Behavior of Thiocyanates on Ignition

The thiocyanates of the alkalies melt readily, and are colored successively
yellow, brown, green, and finally blue, becoming white again on cooling.

The thiocyanates of the heavy metals are decomposed into sulfide, splitting
off carbon disulfide, dicyanogen, and nitrogen. Thus cuprous thiocyanate
is decomposed according to the following equation:

4Cu,(CNS), =4Cu,S+2C8, f+3(CN). T +N: T.

The mercuric thiocyanates swell tremendously on being heated (Pharaoh’s
serpents),

COBALTICYANIC ACID, H;[Co(CN).¢]

The free acid can be obtained by suspending the lead or copper salt in
water and saturating the water with hydrogen sulfide gas; the lead or copper
cobalticyanide is changed into less soluble sulfide and, by filtering, an aqueous
solution of cobalticyanic acid is obtained from which the solid acid deposits
upon evaporation. The free acid may be prepared also by treating the potas-
sium salt with nitric acid, evaporating to dryness on the water bath and extract-
ing the acid with alcohol. After evaporating off the alcohol, needle-shaped,
hygroscopic crystals of H;[Co(CN).] are obtained.

Cobalticyanic acid is extremely stable. It is not decomposed by boiling
with concentrated hydrochloric or nitric acid, by chlorine or by boiling with
HgO. By heating with concentrated sulfuric acid, however, it is decomposed
with evolution of carbon monoxide and carbon dioxide:

2H;[Co(CN),]+8H2SO0.+ 13H20 =2CoSO.,+6(NH4)2S80,+11C0 Tf +COz T.

Solubility of Cobalticyanides—The alkali, alkaline-earth, ferric and other
trivalent metal salts are soluble in water. Most of the other salts witl:
bivalent metals and with the heavy metals are insoluble in water and in acids

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid. No reaction. }

2. Concentrated Sulfuric Acid decomposes all of the salts with evolution
of CO and CO, and formation of blue, anhydrous cobaltous salt.

3. Silver Nitrate produces a white precipitate insoluble in nitric acid but
soluble in ammonia.

4. Lead Acetate, Ferric Chloride and Mercuric Chloride produce no
precipitates. ,

5. Cobaltous Nitrate produces a pink precipitate insoluble in nitric acid
but soluble in ammonia.

6. Nickel Sulfate gives a blue precipitate insoluble in nitric acid but soluble
in ammonia.

7. Copper Sulfate produces a light-blue precipitate insoluble in nitric acid
but forming a blue solution with ammonia.

8. Cadmium Sulfate and Zinc Sulfate give white precipitates insoluble
in nitric acid but soluble in ammonia.

9. Ferrous Sulfate produces a white precipitate insoluble in nitric acid.

10. Mercuric Nitrate produces a white, voluminous pisenitars insoluble
in nitric acid.

REACTIONS IN THE DRY WAY

All the salts of eobalticyanic acid are decomposed by ignition, leaving
cyanide and cobalt carbide behind. The alkali and alkaline-earth salts color
the borax bead blue in both the oxidizing and reducing flames.

328

GROUP II

Silver Nitrate produces a precipitate soluble in nitric acid.
Barium Chloride causes no precipitation.

NITROUS ACID, HNO:

Occurrence.—Nitrous acid never occurs free in nature except in
the form of its salts, the nitrites. It is found in the air, as ammo-
nium nitrite and in many soils and waters, particularly in those which
are contaminated with ammonia or decaying substances.

Ammonia is oxidized by the action of microérganisms (monas
-nitrificans) to nitrous acid, which combines with more ammonia to
form ammonium nitrite.

Preparation of Nitrous Acid and its Salts.—Nitrous acid is formed
by the gentle reduction of nitric acid. If zine is allowed to act upon
dilute nitric acid for a short time, the latter is reduced to nitrous acid,

HNO3+He2=H20+HNOsz,

but the reduction can easily go a little farther, forming NO, N20,
and Ne; while by long-contained action of the zinc, hydroxylamine,
NH20H, and even ammonia are formed.

If nitric acid of sp.gr. 1.3 is heated with arsenious acid, starch,
etc., a mixture of nitric oxide and nitrogen peroxide is obtained, which,
on cooling to —21° C., condenses to a bluish-green- liquid, N2Oz, the
anhydride of nitrous acid.

If the anhydride is treated with ice-cold water, a bluish-green
liquid is obtained, which contains nitrous acid, but always in com-
pany with nitric acid; for N2O3 unites with water, forming nitrous
and nitric acids, and nitric oxide:

2N203+H20 =HNOs+HNO2+2NO0 fT.

At a higher temperature nitrous acid is gradually changed into
nitric acid:

3HNO2=HNO3+2NO Tf +H20.

Pure nitrous acid, therefore, is not known.
If the above mixture of nitric oxide and nitrogen per oxide is con-
329

330 REACTIONS OF THE ACID CONSTITUENTS

ducted into concentrated sulfuric acid, the two gases are readily
absorbed, forming nitrosyl sulfuric acid:

2H2804+NO+NO2 — H20+2H(NO)SOs.

This solution of nitrosyl sulfuric acid in sulfuric acid is sometimes
called ‘‘nitrose.”’
If the solution is added to cold water, sulfuric and nitrous acids

are formed:
H(NO)SO4+H20 — HNO2+H2S04.

A solution of nitrosyl sulfuric acid can be kept indefinitely, so
that it is a convenient reagent for the immediate production of nitrous
acid at any time.

The salts of nitrous acid, the nitrites, are much more stable than
the free acid, and may be obtained by the ignition of nitrates:

2NaNOz = 2NaNOoe +02 T =

Nitrites prepared in this way always contain some oxide and some:
nitrate as impurity.* In order to obtain a pure nitrite, silver nitrite
is treated with the calculated amount of a metallic chloride:

AgNOo+NaCl=AgCl+NaNOsg.

The soluble nitrite can be separated from the insoluble silver
chloride by filtration. 7

Solubility of Nitrites—AlIl nitrites are soluble in water; but silver
nitrite and potassium cobaltic nitrite are difficultly soluble.

REACTIONS IN THE WET WAY

As all nitrites are soluble in water, the reactions which serve for
the detection of this acid cannot be those of precipitation, but rather
those in which a change of color takes place, owing to an oxidation or
reduction.

Nitrous acid sometimes acts as an oxidizing agent, and sometimes
as a reducing agent.

1. Dilute Sulfuric Acid decomposes all nitrites in the cold, setting
free brown vapors:

(a) NaNOo+He2S04 = NaHSO4+ HNO>:;
(b) 3HNOe=HNOs+2NO 7 +H20;
(c) 2NO+02(air) =2NO2 T .

*If the nitrate is heated with a metal, e.g., lead, the reduction takes place at
a lower temperature and is almost quantitative.

NITROUS ACID 331

2. Concentrated Sulfuric Acid reacts exactly the same, but much
more violently.

3. Silver Nitrate precipitates from nitrite solutions crystals of
silver nitrite in the form of needles, which are slightly soluble in cold
water (300 cc. of water dissolve 1 gm. of silver nitrite at, the ordinary
temperature). In boiling water, silver nitrite is considerably more
soluble.

4. Cobalt Salts produce (with an excess of potassium nitrite and
acetic acid) a yellow crystalline precipitate of potassium cobaltic
nitrite (cf. p. 182).

5. Indigo Solution is completely decolorized by warming with
nitrous acid.

6. Hydriodic Acid is oxidized by nitrous acid with separation of
iodine:

2I-+2N0.2-+4Ht — 2H20+2NO tT +To.

If, therefore, a nitrite is added to a solution of potassium iodide
and the solution is acidified with sulfuric or acetic acid, the solution
becomes yellow, owing to the separation of iodine. If the solution
is now shaken with chloroform or carbon disulfide, the latter will
be colored reddish-violet. Or, if a little starch paste is added, it will
be colored blue by the iodine.

As the above equation shows, hydrogen ions are required in the reaction.
If considerable alkali acetate is present, there is no separation of iodine on the
addition of acetic acid, but if a few drops of a strong mineral acid are added,
- jodine is at once set free. This is a good illustration of the mass-action principle
(p. 18) and common ion effect (p. 45).

This exceedingly delicate reaction is also caused by the action of a great
many other oxidizing agents; and it can only be used for the detection of nitrous
acid when it is known that all such oxidizing substances are absent.

As ferric salts also cause liberation of iodine (cf. p. 37) it is evident that
nitrous acid cannot be tested for by the above test in the presence of a ferric
salt. If, however, a large excess of sodium phosphate is added to the solution,
together with a little potassium iodide and some sulfuric acid, no iodine will be
liberated by the ferric salt and the presence of a trace of nitrous acid may
be detected. The ferric ions are converted into very slightly ionized ferric
acid phosphate in this test and the concentration of the ferric ions is made so
small that there is no appreciable reaction with the iodide ions. Artmann,*
who suggested this procedure, adds 8 gms. NazHPQ,-12H,0, 0.2 gm. KI, 5
ec. 4-normal H:SO, and a little starch solution to 100 cc. of a solution of
the ferric salt which is to be tested for nitrite. If as much as 0.3 mg. of NOs
is present, the intense blue color will be obtained immediately.

7. Ferrous Salts are oxidized to ferric salts, with evolution of
nitric oxide:

Fett++N02-+2Ht > Fe**+*+N0O fT +H20.
* Chem., Zlg., 1918, 501.

332 REACTIONS OF THE ACID CONSTITUENTS

The nitric oxide dissolves, in the cold, in the excess of ferrous salt,
forming a brown compound of a varying composition: (FeSO4)2(NO)p.

To obtain this compound, add a little acid to a concentrated solution of
ferrous sulfate and carefully pour the solution to be tested on top. At the zone
of contact between the two solutions the dark-brown coloration will be apparent.

Nitric acid gives the same reaction, but only on the addition of concentrated
sulfuric acid.

8. Potassium Permanganate.—If nitrous acid is added to a warm
acid solution (about 40° C.) of potassium permanganate, the latter
will become decolorized, owing to the oxidation of the former to nitric
acid:

2Mn04-+5NO2-+6Ht — 2Mn*t+5N0O3_+3H20.

In this reaction nitrous acid acts as a reducing agent.

9. Detection of Small Amounts of Nitrous Acid by the Peter Griess
Method.*—To detect the small amounts of nitrous acid which may be present
in drinking-water, of the above reactions only that of potassium iodide and
starch is delicate enough. But as hydrogen peroxide and ferric salts are also
likely to be present, both of which cause the separation of iodine from an acid
solution of potassium iodide, it is evident that dependence upon this reaction
alone would often lead to error.

Consequently, in order to detect the presence of traces of nitrous acid we
make use of a reaction which was first proposed by Peter Griess, and which is
caused by nitrous acid only. It depends upon the formation of an intensely
colored azo dyestuff.

Peter Griess used as his reagent phenylene diamine, whereby a yellow dye,
Bismarck brown, is formed. Recently, according to the suggestion of Ilosvay
v. Ilosva,} an acetic acid solution of sulfanilic acid and of a-naphthylamine is
used instead. According to Lunge,t it is best to mix the solutions of the last
two reagents.

The reagent is prepared as follows:

1. Dissolve 0.5 gm. of sulfanilic acid in 150 ce. of 2-normal acetic acid.

2. Boil 0.2 gm. of solid a-naphthylamine with 20 cc. of water, pour off the
colorless solution from the bluish-violet residue, and add to the colorless solu-
tion 150 cc. of 2-normal acetic acid.

Mix the two solutions. The mixture keeps well only when kept in a dark
place. It turns reddish if exposed to the light, and cannot then be decolorized
by shaking with zinc.

Procedure.—Treat about 50 cc. of the water with 2 cc. of the above reagent,

and allow it to stand five or ten minutes; it will be colored a distinct red if a
trace of nitrous acid is present.

10. Diphenylamine, dissolved in concentrated sulfuric acid, is

* Ber., 12, (1879) 427.
¢ Bull. chim. [3] 2, 317.
¢ Z. angew. Chemie, 1889, Heft 23.

NITROUS ACID 333

colored intensely blue by nitrous acid. Nitric acid and many other
oxidizing substances, such as selenic acid, chloric acid, ferric chloride,
etc., will give the same reaction (cf. Nitric Acid).
11. Brucine dissolved in concentrated sulfuric acid (according
to G. Lunge and A. Lwoff *) gives no reddish coloration when treated
with nitrosyl sulfuric acid.

Dry, repeatedly-recrystallized silver nitrite, containing 70.05 per cent.
silver (theory 70.09), did give with brucine (cf. p. 395), in an atmosphere of
carbon dioxide, a weak but nevertheless distinct test for nitric acid. probably
due to the presence of traces of nitrate remaining in the silver nitrite. On
dissolving 15 mgm. of this same nitrite in water, adding an equivalent amount
of sodium chloride and diluting to one liter, a solution of sodium nitrite was
obtained, of which one ce. added drop by drop with constant stirring to about
4 cc. of concentrated sulfuric acid yielded a solution of nitrosyl sulfuric acid
which gave no sign of a red coloration with a drop of brucine reagent. The test
was immediately obtained, however, on adding a trace of nitric acid to this
solution.
Brucine, therefore, is a reagent by which nitric acid can be detected in the
presence of nitrous acid.

12. Urea acts with nitrous acid with evolution of nitrogen and
carbon dioxide:

CO(NH2)2+2HNO2 > 3H20+C02 T +2N2 7.

In this reaction the nitrogen of the urea, with its negative valence
of three, is oxidized by the nitrogen of nitrous acid, which has a positive
valence of three. In alkaline solutions free halogens will react with
urea, but in acid solutions the reaction with nitrous acid is sensitive
and characteristic. ;

13. Ammonium chloride, on being added to a boiling solution of a
nitrite in dilute acetic acid, causes evolution of nitrogen gas (cf. p. 99):

NH,t+N0O2°— 2H20+Noe2 T ;

HYDROGEN SULFIDE (HYDROSULFURIC ACID), H2S

Occurrence and Preparation.—Hydrogen sulfide is found in volcanic
regions, in many mineral waters (the so-called ‘ sulfur” waters),
and, in general, wherever substances containing sulfur are subject to
decay; or when they come in contact with decaying substances. Sul-
fates are easily changed into sulfides by the action of microérganisms
which are present in the air; and this is the reason why many mineral
waters containing sulfates smell of hydrogen sulfide after standing some
time in a corked flask. If, however, the flask and the cork are sterilized,
the water can be kept indefinitely. The formation of hydrogen sul-
fide from sulfates takes place as follows:

By means of carbonaceous matter (dust, etc.) the sulfates are
reduced with the aid of microérganisms, at first to sulfides.

Na2SO4+Ce2=2CO02+Naz8S,
which are then decomposed by carbonic acid:
NaeS a HeCO3 = NaeCO3 +H.S.

Just as hydrogen sulfide may be made from sodium sulfate by
the action of organic matter in a corked flask, so in nature the same
process brings about the presence of hydrogen sulfide in many mineral
waters.

For laboratory purposes, hydrogen. sulfide is. similarly prepared
by the action of dilute sulfuric or hydrochloric acid upon a sulfide
(usually iron sulfide, FeS, on account of its cheapness and stability).

Properties—Hydrogen sulfide is a colorless gas, with an odor
like that of rotten eggs; it is absorbed by water at the ordinary tem-
peratures (one volume water absorbs two to.three times its own
volume). The saturated solution at 25° is approximately 0.1.molal=
0.2 normal. The higher the temperature, the less the solubility, as
with all gases. For the primary ionization: H2S @H++HS-, the
value of the ionization constant

(2g bod: bp z
(Hs) =k, is 0.911077.
For the secondary ionization, HS~ - H++S-, the value of the constant
[H+] x[S=]

[HS] =ke is 1.210-5,
334

HYDROGEN SULFIDE 335

In the saturated solution at 25°, the concentration of the hydrogen
ion is about 0.9X10~ molar equivalents per liter and of the simple
sulfide ion it is only 1.2X10°'. According to the table on p. 56,
therefore, a saturated solution of hydrogen sulfide should react acid
to phenolphthalein and to azolitmin but not to methyl orange.

The solution of hydrogen sulfide becomes turbid on standing in
the air as a result of its oxidation by atmospheric oxygen:

2H28 + O02=2H20+S8o.

Hydrogen sulfide burns in the air with a bluish flame to water
and sulfur dioxide:
2H2S +302 =2H20+2S02.

The salts of hydrosulfuric acid are called sulfides.

Solubility cf Sulfides—The sulfides of the alkalies and the hydro-
and polysulfides of the alkaline earths are soluble in water. The
monosulfides of the alkaline earths, particularly calcium sulfide, CaS,
are difficultly soluble in water, but they are gradually changed from
contact with water into soluble hydrosulfides:

2CaS+2H20 — Ca(OH)2+Ca(SH)o.

The remaining sulfides are insoluble in water. Of these latter
FeS, MnS, and ZnS are decomposed by dilute hydrochloric acid with
evolution of hydrogen sulfide; others require concentrated hydro-
chloric acid, e.g., SbeS3, SnSe, PbS, NiS, CoS, CdS; while the remain-
ing are insoluble in concentrated hydrochloric acid, but are all soluble
in aqua regia with separation of sulfur.

REACTIONS IN THE WET WAY

Free hydrogen sulfide, as has been stated, is a very weak acid, being even
weaker than carbonic acid. The soluble neutral salts R.S are decomposed into
metal and sulfur ions,

RS = 2Rt+S-,
but under the influence of water, some of the bivalent sulfur ions are changed
to univalent HS ions,
HOH+S~- — HS-+0H;
some of the bivalent sulfur ions remain in solution, and, in fact, more in con-
centrated solutions than in dilute ones.

As, therefore, an aqueous solution of a sulfide contains both S ions and SH
ions, while the solutions of the free acid contain chiefly SH ions, it is plain why
in many reactions the former react in a somewhat different way from the latter.

1. Dilute Sulfuric Acid decomposes all soluble, and some insoluble,
sulfides, with evolution of hydrogen sulfide.

336 REACTIONS OF THE ACID CONSTITUENTS

2. Concentrated Sulfuric Acid decomposes all sulfides, on warm-
ing, with evolution of sulfur dioxide and deposition of sulfur:

NaoS+2H2804 = Na2gSO4+2H20+802 fT +8.

But even the sulfur goes over into sulfur dioxide after being heated
with the sulfuric acid for some time:

2H2S04+S =2H20+3S802 fF.

8. Silver Nitrate produces, in solutions of hydrogen sulfide and
of soluble sulfides, a black precipitate of silver sulfide,

2A0++HS > AgsS+2H*,

insoluble in cold nitric acid, in which, however, it dissolves on warm-
ing (ef. p. 181).

4. Barium Chloride causes no precipitation.

5. Lead Salts (best a solution containing an excess of alkali) pro-
duce a black precipitate of lead sulfide. . All sulfides which are decom-
posed by hydrochloric acid evolve hydrogen sulfide, which, on coming
in contact with a piece of filter-paper moistened with an alkaline lead
solution, colors the latter black. An insoluble sulfide (pyrite, arsenic
sulfide, mercuric sulfide, etc.), evolves hydrogen sulfide with hydro-
chloric acid and nascent hydrogen.

“\

To test an insoluble sulfide, such as the mineral pyrite, place a little finely
granulated tin in a test tube, cover it with 6-normal hydrochloric acid and heat
gently. Hold a piece of filter paper which has been moistened with lead acetate
solution and a little ammonium hydroxide, over the escaping vapors. If it
blackens, the tin itself contains a little sulfide and cannot be used for the most |
delicate test. Usually, however, the blackening with tin alone is so slight that
an allowance can be made for it. Now add a little of the substance to be tested
and a little more tin. A trace of sulfide will cause blackening of the lead
acetate paper.

6. Sodium Nitroprusside, Nag{Fe(CN)s(NO)]-2H20, is colored
reddish-violet by 5 ions, but not by SH ions. Consequently hydrogen
sulfide itself does not give this reaction, except upon the addition
of caustic ‘alkali. The reaction is very sensitive, but not so delicate
as the one with an alkaline solution of a lead salt.

7. Methylene Blue.—This reaction (which was recommended by
Emil Fisher *) is the most sensitive of all reactions for detecting the
presence of hydrogen sulfide. It is particularly suited for detecting
the presence of traces of. hydrogen sulfide in mineral waters, evep —
when all other tests give negative results.

* Ber., 16, 2234.

HYDROGEN SULFIDE 337

Treat the solution to be tested for hydrogen sulfide with one tenth of its
volume of concentrated hydrochloric acid, add a little dimethylparaphenylen-
diamine sulfate, NH.-CsH,-N(CH;)2-HSO,, from the point of a knife-blade,
stir it into the liquid, and when it has dissolved, add one or two drops of a
dilute ferric chloride solution.

The formation of methylene blue may be expressed by the following equation:

2[NH, , C.H, ’ N(CHs): , HCl] +6Fet + F4e§=

Jf Otln —N (CHs)2
S +6Fet ++NH,++4Ht+Cr.

—N
NH; —==N (CHs)2

If only 0.02 mgm. of hydrogen sulfide is present in a liter, the blue color is
distinctly apparent after half an hour’s standing, while the above tests would
give negative results.

If too little hydrochloric acid is present, a red coloration is obtained; this
is caused by the action of ferric chloride upon a faintly acid solution of di-
methylphenylendiamine. If considerable hydrochloric acid is present, the red
coloration does not appear.

8. Oxidizing Agents, such as the halogens, nitric acid, chromates,
permanganates, ferric salts, etc., decompose hydrogen sulfide with
separation of sulfur.

9. Metallic Silver is blackened by both free hydrogen sulfide and
soluble sulfides:

2Ag+H2S+0 (air) =H20+Ag.S;
2Ag+Na2S+H20+0 (air) =2Na0H+Ag.S.

If oxygen and water are not present, the above reactions will not
take place. A piece of bright silver suspended for fourteen hours in
a sulfur spring showed no sign of darkening until it had been exposed
~ to the air for a short time. )

Absolutely dry hydrogen sulfide, in the presence of absolutely
dry oxygen, acts upon silver at ordinary temperatures only very slowly;
it acts instantly if a trace of water is present.

To detect the presence of sulfur in insoluble sulfides, fuse with a little
caustic soda (on the cover of a porcelain crucible) to form soluble sodium sulfide:

NiS+2Na0H =H.,0+Ni0O+Na8.

Some sulfate is always formed by this treatment; but the aqueous solution
of the melt will always contain enough alkali sulfide for any of the above tests.

338 - REACTIONS CF THE ACID CONSTITUENTS

Behavior of the Sulfides on Ignition

Most sulfides remain unchanged when heated out of contact with
the air; arsenic and mercuric sulfides sublime. ’

The polysulfides lose sulfur, which sublimes. The sulfides of gold
and platinum lose sulfur, leaving the metal behind. All sulfides when
heated in the air give off sulfur dioxide, which can be recognized by
its odor.

The Detection of Sulfur in Non-Electrolytes is usually effected by
heating the substance in a glass tube with metallic sodium (ef. p. 291), and
testing the aqueous extract of the melt with sodium nitroprusside; or, the resi-
due in the tube may be treated with dilute hydrochloric acid and the escaping
gas tested with lead acetate paper for hydrogen sulfide.

The following method of testing for sulfur is very certain. It
depends upon the conversion of any sulfur present into the sulfate
ion which is tested for barium chloride in hydrochloric acid solution.
(Cf. p. 404). The best way of converting the sulfur into sulfuric
acid is to heat with concentrated nitric acid in a sealed tube. | (Carius
method, see Vol. II) or, in the case of difficultly volatile substances
low in sulfur, by fusion with sodium peroxide in a nickel crucible.
Since, however, this last oxidation often takes place with explosive
violence, it is best not to use pure sodium peroxide, but to mix it with
sodium-potassium carbonate. Mix the substance (from 0.1 to 5
gms. according to the sulfur content) with ten times as much sodium-
potassium carbonate and three times as much sodium peroxide and
heat in a nickel crucible, with the crucible inserted in a disk of
asbestos board to keep the flame of the gas away from the contents;
if this precaution is not taken a little sulfate is obtained from the
sulfur in the gas. After cooling the melt, dissolve it in water, filter,
acidify with hydrochloric acid and test the filtrate for sulfuric acid
with barium chloride, | .

SULFUR, S

M. Pt. =113.5°—119.5°; B. Pt. 444.5°.

Occurrence.—Sulfur is found native in voleanic regions in the form of ortho-
rhombic pyramids, and in the neighborhood of sulfur waters, being formed from
the oxidation of some of the hydrogen sulfide which escapes.

Preparation and Properties.—Like the halogens, sulfur is formed by the
oxidation of its hydrogen compound:

By heating polysulfides or the sulfides of shes noble metals (gold and plati-
num), sulfur is also obtained.

Sulfur exists in three allotropic modifications:

1. As Orthorhombic Sulfur, with a melting-point of 113.5° C., obtained by
crystallization from solutions below 95°.

2. As Monoclinic Sulfur, with a melting-point of 119.5° C., obtained by the
solidification of molten sulfur.

3. As Amorphous Sulfur, obtained by quickly cooling the molten sulfur after
it has become viscous by heating to 250° C., or after it has become a thin a
after heating to a higher temperature.

Monoclinic sulfur changes gradually into orthohombic octahedrons; or,
in other words, the unsymmetrical form changes into the symmetrical form.
This is a general phenomenon:

If a substance exists in two or more crystallographic forms, the more symmetrical
form is always the more stable, and the less symmetrical always has a tendency to
go over into the more symmetrical form. Thus the unsymmetrical, orthorhombic,
yellow mercuric iodide is changed, by rubbing with a glass rod, into symmetrical,
tetragonal, red mercuric iodide (cf. p. 196); and, similarly, the orthorhombic
form of calcium carbonate, aragonite, changes into hexagonal calcite.

Both of the crystalline modifications of sulfur are soluble in carbon disul-
fide; and, by evaporating the solution, the sulfur always recrystallizes in the
form of octahedrons. Amorphous sulfur is insoluble in carbon disulfide.

Commercial “‘ flowers of sulfur” is a mixture of crystalline and amorphous
sulfur, and therefore is only partly soluble in carbon disulfide.

Sulfur burns in the air to sulfur dioxide, and, in the presence of “ contact
substances,”’ such as platinum, oxide of iron, chromic oxide, etc., it is burned
to sulfur trioxide also. Consequently the gas from pyrites burners always
contains a mixture of the two gases.

Sulfur is insoluble in water, but soluble in hot caustic alkali, forming a,
thiosulfate and a sulfide:

6Na0OH +48 = 3H,0 +. Na.S.0; 4+ 2Na.S.
By further action of sulfur the Na2S is changed into NaS:, NasSs, ete.

This reaction is entirely analogous to the formation of hypochlorite and
chloride by the action of chlorine upon cold dilute caustic alkali:

2NaOH+Cl, = H.0+Na0Cl+NacCl.
339

ole in alkali sulfites, forming a t

. f ‘ .

seta -—NaSOstS=NaS.0.,

ACETIC ACID , HC2H302

Acetic acid is found in the sap of many plants, partly free and
partly in the form of its potassium or calcium salt.

It is formed by the dry distillation of wood or by the oxidation
of alcohol. |

Anhydrous acetic acid (glacial acetic acid) solidifies below +16°
C., forming colorless, glistening plates. It has a penetrating odor,
similar to that of sulfur dioxide, and is miscible with water, alcohol,
and ether in every proportion. It boils at 118° C.

The aqueous solution reacts acid. It is monobasic acid, and
its salts, the acetates, are as a rule readily soluble in water; the silver
salt is difficultly soluble.

The most important commercial salts of this acid are sodium
acetate and lead acetate (sugar of lead), Pb(C2H302)2+3H20.

Neutral lead acetate dissolves lead oxide with the formation of
soluble basic salts:

—_OH

Pb(C2H302)2+PbO+H20=2Pb_ G57, 3
—O—Pb—CsH30>

Pb(C2Hs02)2+2PbO=Pb_ 4: bh 6.7.05"

The solutions of the soluble basic lead acetates, as. well as that of neutral
lead acetate, yield precipitates of lead carbonate when acted upon by carbon
dioxide. For this reason turbid solutions are often obtained when these salts
are dissolved in distilled water, because the latter frequently contains carbonic
acid. If a drop of acetic acid is added to the turbid solution the precipitate
disappears at once.

REACTIONS IN THE WET WAY

Use a solution of sodium acetate for the following reactions:

1. Dilute Sulfuric Acid sets acetic acid free from its salts; it is
quite volatile and can be recognized by its odor.

2. Concentrated Sulfuric Acid also sets acetic acid free. If alcohol
is added at the same time and the mixture warmed, ethyl acetate

is formed, )
HC2H302+C2Hs50H = H20+Ce2Hs-CoH302,

Ethylacetate

which can be recognized by its pleasant, fruity odor.
341 °

342 REACTIONS OF THE ACID CONSTITUENTS

3. Silver Nitrate produces, in fairly concentrated solutions, a —
white crystalline precipitate of silver acetate (100 parts of water
dissolve 1.04 parts at 20° C. and 2.52 parts at 80° C.).

4, Ferric Chloride colors neutral solutions dark brown, and by
boiling the diluted solution, all the iron is precipitated as basic acetate
(ef. p. 150).

REACTIONS IN THE DRY WAY

All acetates are decomposed on ignition, leaving behind either
the carbonate, oxide, or the metal itself, and with the evolution of
combustible vapors and gases.

The acetates of the alkalies are decomposed into carbonate and
acetone:

2NaC2H302 = NagCO3+(CHs3)2CO.

The acetates of the alkaline earths always leave the metal in the
form of its oxide, while the acetates of the noble metals leave a resi-
due of the metal itself.

Cacodyl Reaction.—If a dry acetate (best an alkali acetate) is
heated with arsenic trioxide, a very repulsive-smelling and extremely
poisonous gas is formed, called cacodyl oxide:

4NaCoH302 + AseOz et 2Na2CO3+[(CH3)2As]20+2CO02 T °

In spite of the sensitiveness of this test, it cannot always be
relied on, for many other organic acids, such as butyric and valeri-
anic acids, give similar reactions.

Cyanic AcID, HCNO

This very unstable acid is obtained by heating its polymer, cyan-
uric acid, (HCNQ)3; it is a colorless liquid with a very penetrating,
disagreeable odor, which immediately decomposes in aqueous solution,
forming ammonium bicarbonate:

HCNO+2H20 — HCO2NH2+H20 — NH4HCOs.
Carbamic acid -

The salts of cyanic acid, the cyanates, are much more stable than
the free acid, and may be obtained by the oxidation of cyanides.

By simply fusing potassium cyanide in the air, perceptible quan-
tities of potassium cyanate are formed. If, however, potassium
cyanide is heated with oxidizing substances, or those which can be
readily reduced, it is easy to change the cyanide completely over
to cyanate. ‘The cyanates of the alkalies are stable in the dry state,
but take on moisture from the air and are gradually changed into
alkali bicarbonate and ammonia:

KCNO+2H20 =KHCO3+NHs ¢.

Solubility of Cyanates—The cyanates of the alkalies and alkaline
earths are soluble in water. Silver, mercurous, lead, and copper
cyanates are insoluble in water. All cyanates are soluble in nitric
acid.

REACTIONS IN THE WET WAY

Use a freshly prepared, cold solution of potassium cyanate for
these reactions.

1. Dilute Sulfuric Acid immediately sets cyanic acid free, which
decomposes into ammonium salt and carbon dioxide:

CNO-+H*t — HCNO; HCNO+2H20 — NH4*+HCOs;";
HCO3-+H* — H20+C02 fT.

Consequently a strong evolution of carbon dioxide takes place on
adding the sulfuric acid. The carbon dioxide always contains small
amounts of undecomposed cyanic acid, which is recognizable by its
very penetrating odor. The solution will contain ammonium sulfate;
if it is warmed with caustic soda, ammonia will be given off.

343

344 REACTIONS OF THE ACID CONSTITUENTS

2. Concentrated Sulfuric Acid reacts similarly.
3. Silver Nitrate precipitates white, curdy silver cyanate,

KCNO+AgNO3=KNO3+AgCNO,

soluble in ammonia and in nitric acid (difference from silver cyanide).

4. Barium Chloride produces no precipitation.

5. Cobalt Acetate is colored azure blue by a solution of potassium
cyanate. The blue potassium cobaltocyanate, Ke[Co(CNO)a4], dis-
covered by Blomstrand,* ‘is formed by this reaction, and is obtained
in the form of tetragonal crystals of a dark azure-blue color. :

This blue compound dissolves in water with a blue color. If, however,
it is subjected to the action of considerable water, the color disappears, the
double salt being dissociated into its components:

K,Co(CNO), = Co(CNO).+2KCNO.

If more potassium cyanate is added to the solution, which has become
colorless, the blue color reappears. The same result is reached by adding
alcohol.

Almost all commercial potassium cyanide contains some cyanate.

In order to detect the presence of cyanate in the commercial salt, the cyan~
ide must first be expelled, for the cobalt test does not take place in the presence
of cyanide. |

According to E. A. Schneider,} the test is made as follows:

Dissolve 3 to 5 gms. of the potassium cyanide in 30 to 50 ec. of cold water
and pass carbon dioxide into the solution for 60 to 90 minutes; the hydrocyanic
acid is expelled, and potassium bicarbonate carbonate is formed, while the
potassium cyanate is not affected perceptibly:

KCN+H:CO;=HCN f +KHCO;.

Take 1 ce. of the solution, add 25 cc. of absolute alcohol (to precipitate
potassium bicarbonate) and filter. Treat the alcoholic filtrate with a few
drops of acetic acid and then add a few drops of alcoholic cobalt acetate solution.

If the original cyanide contained 0.5 per cent of potassium cyanate, the blue
color can be distinctly recognized, but alkali thiocyanates give the same reaction.

* Journal fir praktische Chemie [2], 3, 206.
T Ber., 1895, p. 1540.

HYPOPHOSPHOROUS ACID, H3POz2

Hypophosphorous acid is obtained by the decomposition of its
barium salt with sulfuric acid, or of its calcium salt with oxalic acid.
The acid is monobasic, only one of the hydrogen atoms being replace-
able by metals. The salts of hypophosphorous acid are obtained
by boiling phosphorus with dilute alkali, whereby phosphine is given
off :

2P4+3Ba(OH)2+6H20 =3Ba(H2PO2)2+2PHs3 f ;

P1+3KOH+3H20 =3K(H2PO2)+PH3 T ;

The phosphine thus obtained is spontaneously combustible owing
to the presence of small quantities of liquid phosphine, P2H4; it is,
however, mixed with considerable hydrogen because the alkali reacts
upon the alkali hypophosphite with evolution of hydrogen. See No.
6, below.

- Solubility of Pr spenkoucied: —All hypophosphites are soluble in
water.

REACTIONS IN THE WET WAY

1. Dilute , Sulfuric Acid.—No reaction.

2. Concentrated Sulfuric Acid reacts with hivnaplicenhtea only
on warming, and is reduced to sulfur dioxide, which can be recog-
nized by its odor.

3. Silver Nitrate is reduced to metallic silver, sometimes with and
sometimes without the evolution of hydrogen, according to the relative
amounts of the substances reacting:

2H2PO2 +2Ag*+4H20 — 2H3P04+2Ag+3Hpe fT ,
HePO2 +4Agt+2H20 — H3P04+4Ag+3H". ©

4. Barium Chloride causes no precipitation.
5. Copper,* Mercury, and Gold Salts are reduced to metal.
6. Concentrated Caustic Potash.—By boiling with concentrated

* With copper the reduction may go so far that copper hydride is formed. Cf.
" Wirz, Compt. rend., 18, 102.
345

346 REACTIONS OF THE ACID CONSTITUENTS
caustic alkali, the hypophosphites are oxidized, with evolution of
hydrogen, to phosphates:
H2PO2°+20H™ — POe+2He J.
7. Nascent Hydrogen (zinc and dilute sulfuric acid) reduces hypo-
phosphorous acid to phosphine (see Phosphorous Acid).
REACTIONS IN THE DRY WAY
By ignition phosphate and phosphine are obtained:
2H3PO02 — H3P04+PHs f ;
4NaH2PO2 — NasP207+H20 f +2PHs 7 ;
2Ca(H2PO2)2 — CazP207+H20 7 +2PH3 7.

GROUP III

Silver Nitrate produces a white precipitate, soluble in nitrie acid.
Barium Chloride does the same.

SULFUROUS ACID, H2SO;

— Occurrence and Preparation—Sulfur dioxide, the anhydride of
sulfurous acid, is found in the exhalations of active volcanoes, and
is formed by the combustion of sulfur or sulfides in the air.

S+02=S0O2 fT,
4FeSo+1102 = 2Fe203+8S02,

or by the reduction of sulfuric acid on heating with sulfur, sulfides,
carbon, organic substances, and metals:

2H2804+8 =2H20+3S02 T ;
2H2804+C =2H20+CO2 fT +2802 i gr
2H2804+ Cu =2H20+CuS04+SO02 1 OF

Mercury, silver, tin, etc., act the same as copper.
Sulfur dioxide is also formed by the decomposition of sulfites and
thiosulfates with stronger acids:

NazSOs-+H2804=Na2S04+H20+S80z f ;
NaeSe03+ HeS04 = NaeSO4+8+He0+S802 tT ;

Sulfurous acid may be prepared for laboratory purposes by placing
a concentrated solution of sodium bisulfite in a flask and allowing
-concentrated sulfuric acid to drop upon it. A steady stream of sulfur
dioxide will be evolved without warming. . :

Sulfur dioxide is a colorless gas, having the penetrating odor
peculiar to burning sulfur, and is readily soluble in water and alcohol:
one vol. water at 15° C. dissolves 45.36 vol. SQ2; one vol. alcohol at
15° C. dissolves 116 vol. SOz.

The aqueous solution contains sulfurous acid, H2SO3. The acid
cannot be isolated, as it decomposes on evaporation into water and
sulfur dioxide; consequently the free acid is known only in aqueous

347

348 REACTIONS OF THE ACID CONSTITUENTS

solution. By neutralization of this solution with alkali hydroxides
or carbonates, the comparatively stable sulfites are obtained. Jn
solution, sulfites are gradually oxidized to sulfates by dissolved oxygen.

Solubility of Sulfites——The sulfites of the alkalies are readily
soluble in water; the remaining sulfites are difficultly soluble or in.
soluble in water, but are all soluble in hydrochloric acid.

REACTIONS IN THE WET WAY

1. Dilute Sulfuric Acid evolves sulfur dioxide, in the cold, from
all sulfides, the gas being easily recognized by its odor.

2. Concentrated Sulfuric Acid reacts in the same way, but much
more energetically.

3. Silver Nitrate produces, in neutral solutions of sulfites or in
an aqueous solution of sulfurous acid, a white crystalline precipitate
of silver sulfite,

SO3-+2Agt — AgoSOs,
soluble in an excess of alkali sulfite, forming sodium silver sulfite:
AgeSO3-+ Na2SO3 = 2Nal[AgSOs].
By boiling this solution the silver is precipitated as a gray metal:
2[AgSO3l — SOs" +802 7 +2Ag.

If water containing silver sulfite in suspension is boiled, half the
silver is reduced to metal, while the other half goes into solution as
sulfate:

2Ag2SO3 — =AgeSO4+S8O2 T +Age.

Silver sulfite is soluble in ammonia and in nitric acid.

4, Barium Chloride produces no precipitation in an aqueous solu-
tion of sulfurous acid, but in neutral sulfite solutions white barium >
sulfite is precipitated,

SO3-+Ba** — BaSOs,

readily soluble in cold, dilute nitric acid. By boiling the solution,
barium sulfate is formed slowly and separates out. As sulfites in
aqueous solution gradually change to sulfates, commercial sulfites
usually contain sulfate. In this case the precipitate produced by
barium chloride in neutral solution contains barium sulfate, which
is insoluble in dilute nitric or hydrochloric acid. If the barium sul-
fate is filtered off and the filtrate treated with chlorine or bromine

SULFUROUS ACID 349

water, a white precipitate of barium sulfate is formed, provided a sul-
fite was originally present:

803-+Cle+Ba*t*+H.0 — 2H+ +2Cl-+BaSO,.

5. Strontium and Calcium Salts behave similarly to the barium
salt.

The sulfites of the alkaline earths vary in their solubilities in
sulfurous acid and in water.

Calcium sulfite readily dissolves in an excess of sulfurous acid,
forming calcium bisulfite:

CaSO; + H2SO3 = Ca(HSOs) 2.

On boiling this solution, sulfur dioxide escapes, and calcium sul-
fite is reprecipitated. |

The strontium salt also dissolves in sulfurous acid, but more dif-
ficultly; the barium salt is practically insoluble in sulfurous acid.

SOLUBILITY OF THE ALKALINE-EARTH SULFITES IN WATER

One part calcium sulfite dissolves in 800 parts water at 18° C.
One part strontium sulfite dissolves in 30,000 parts water at 18° C.
One part barium sulfite dissolves in 46,000 parts water at 18° C.

Advantage is taken of the difficult solubility of strontium sulfite
in detecting sulfurous acid in the presence of thiosulfuric acid (which
see). .
6. Lead Salts precipitate white lead sulfite, soluble in cold dilute
nitric acid; but on boiling the solution lead sulfate is precipitated.

7. Sodium Nitroprusside and Zinc Sulfate——If a neutral sulfite
solution is treated with a dilute solution of sodium nitro-prusside, a
faint pink coloration is produced. If, however, considerable zine sulfate
is added, the coloration becomes a distinct red. The reaction is still
more sensitive if a little potassium ferrocyanide is added, a red pre-
cipitate being formed (difference from thiosulfuric acid). This reac-
tion, although very delicate, is not so reliable as the precipitation
with strontium chloride.

Sulfurous acid is a strong reducing agent.

8. Iodine Solutions are decolorized by sulfurous acid:

SO3-+H.0+1. > 2H*+21I-+S80.4".

9. Acid Potassium Permanganate Solutions are also decolorized,
sulfuric and dithionic acids being formed in varying amounts accord-
ing to the temperature and concentration.

350 REACTIONS OF THE ACID CONSTITUENTS |

Under certain conditions the reaction can take place according
to the following equations:

2MnO4~+6S8037-+8Ht — 2Mn*t+4804>+820¢67+4H20.

Under other conditions, however, the sulfurous acid can be com-
pletely oxidized to sulfuric acid. Consequently, sulfurous acid cannot
be determined quantitatively by means of potassium permanganate.

10. Chromic Acid is reduced to green chromic salt:

2CrO4~ + 3503 + 10H* =< 2Crt + Fe 3804->+ 5H20.

11. Mercuric Chloride is unaffected by sulfurous acid at ordinary
temperatures; but, on boiling, it is reduced to mercurous chloride,

2H¢gCle +S8037°+H20 — 2Ht+ 2Cl- +8047 +Hge2Cle :

and on adding more sulfurous acid, the mercurous salt is reduced
to gray metal.

12. Mercurous Nitrate is immediately acted upon by free sul-
furous acid and by alkali sulfite solutions, with the formation of a
black precipitate.

13. Gold Solutions are also reduced.

14. Nascent Hydrogen reduces sulfurous acid to hydrogen sul-
fide, which can be recognized by its odor and by its turning lead
acetate paper black. The reduction is best effected with zinc and dilute
hydrochloric acid.

REACTIONS IN THE DRY WAY

The sulfites of the alkalies, when heated out of contact with the
air, are changed to sulfate and sulfide:

4Na2SO3 ais 4 3Na2S0.,+Na.S.

By heating an alkali sulfite in the closed tube this reaction takes
place, and there is no sublimate of sulfur (difference from thiosul-
fates). If the melt is treated with hydrochloric acid after cooling,
hydrogen sulfide is given off freely.

The remaining sulfites are changed, on being heated out of con-
tact with the air, into sulfur dioxide and oxide or metal:

CaSO3=CaO+S0Osz J ;
2Ag2SO3 = 4Ag+2802 T +Oe T ;

If a sulfite is heated with sodium carbonate on charcoal, sodium
sulfide is formed. If the melt is placed upon a bright silver coin and

SULFUROUS ACID 351

.moistened with water, the silver is blackened, owing to the formation
of black silver sulfide (Hepar reaction) :

2Na2S03+3C =3CO2 T +2Na2S,
and
NaeS+2Ag+H20+0=2Na0H+Ag2S.

This Hepar reaction takes place with all sulfur compounds, and
therefore shows simply the presence of sulfur. The oxygen required
in the above reaction is obtained from the atmosphere.

CARBONIC ACID, H2COz

Like sulfurous acid, pure carbonic acid does not exist; it is known

only in aqueous solution. Its anhydride, COs, is formed by the com-
bustion of carbon and of carbonaceous matter of all kinds, and is found
therefore very widely distributed in nature (in small amounts in the
atmosphere, and in enormous amounts in volcanic regions, streaming
out from fissures in the earth).

Pure air contains 0.35-0.40 per cent of CO:. In dwelling places the amount
increases considerably, owing to breathing and other forms of combustion.
If 3 or 4 per cent is present, as is the case in mines sometimes, breathing
becomes difficult, and the miners’ lamp begins to burn faintly; while when 8
to 10 per cent is present the lamp goes out.

Carbon dioxide occurs also in many mineral waters, and (in the

liquid state) is found enclosed in quartz, feldspar, etc. As carbonate
it exists in enormous quantities as limestone, marble, aragonite, dolo-
mite, etc. Carbon dioxide is a colorless, odorless, slightly acid-tast-
ing gas, with sp.gr. 1.52. Being, therefore, one and one-half times as
heavy as air, it can be poured from one vessel into another. Carbon
dioxide does not support combustion; a burning candle goes out in
air containing 8 to 10 per cent of this gas.*

Carbon dioxide gas is not very soluble in water. At 15° and 760 mm. pres-
sure, 1 liter of water dissolves its own volume of gas. At higher temperatures
it is less soluble and it is easy to expel carbon dioxide from a solution. In the
aqueous solution, the following equilibria exist

CO.+H,0 < H,CO; H++HCO;";
HCO, Ht+Coz".

The ionization constant for the primary ionization of carbonic acid is 0.0.3 and
for the secondary ionization it is 0.0:>7. Aécording to these values, the primary
ionization of carbonic acid takes place to about 0.1 per cent and there is present
in a liter of saturated carbon dioxide solution at 15° only about 0.0006 equiv-
alent of hydrogen ions. The secondary ionization upon which the quantity
of CO;~ ions present depends, takes place only to a negligible extent.

* Carbon dioxide not only fails to support combustion, but it tends to prevent
it; hence its use in fire extinguishers. Being formed by the combustion of car-
bon, the mass-action principle shows that it will tend to stop the reaction much
better than an inert gas.

352

CARBONIC ACID. 353

When'a strong acid is added to the solution, even the primary ionization of
carbonic acid is repressed almost completely. Similarly, when hydrogen ions
are added to a carbonate, carbon dioxide is formed, even with acetic acid, and
the carbon dioxide is easily expelled by heating.

The salts of carbonic acid, the carbonates, are formed:

1. By passing carbon dioxide gas into a solution of a metallic
hydroxide:
: 2Na0OH +COe2 = HeO +NaeCOz ;

Ba(OH)2+CO2=H20+BaCOsz.

2. By the action of carbon dioxide upon cyanides, sulfides, and
borates of the alkalies and alkaline earths.

3. By the ignition of salts of organic acids (cf. p. 80).

An illustration of the preparation of large amounts of carbonate
is the production of potash by burning parts of plants (wood, for
example, or the residue from the manufacture of beet sugar, the latter
being particularly rich in potassium salts).

Solubility of Carbonates.—Of the normal carbonates, only those
of the alkalies are soluble in water; and their aqueous solutions react
alkaline, owing to hydrolytic dinoaipodiiiow:

Na2CO3+H20 @ 2Nat+OH-+HCO3-.

The aqueous solution of the carbonates of the alkalies, therefore,
behaves as if it were a solution of caustic alkali and alkali bicar-
bonate.

Many carbonates dissolve in an excess of carbonic acid, forming
bicarbonates, particularly the alkaline-earth carbonates:

CaCO3+ H2CO3 = Ca(HCOs)e.

By boiling a solution of calcium bicarbonate, the latter is decom-
posed into water and carbon dioxide, and calcium carbonate is repre-
_ cipitated:

Ca(HCQs)2 =H.20+C0O2 T +CaCOsz.

Nearly all samples of drinking-water contain calcium or mag-
nesium bicarbonate; they become turbid, therefore, on boiling (boiler
scale). Dilute, cold mineral acids decompose all carbonates with
effervescence (due to evolution of carbon dioxide gas).

The native carbonates of magnesium and iron (magnesite, siderite,
and dolomite) do not effervesce if a lump of the mineral is treated
with cold dilute mineral acids, but when reduced to a fine powder
they are more readily acted upon; on warming, all carbonates dissolve
readily.

354 REACTIONS OF THE ACID CONSTITUENTS

REACTIONS IN THE WET WAY

1, Dilute Sulfuric Acid decomposes all carbonates with efferves-
cence; except with magnesite, siderite, and dolomite, the reaction
takes place in the cold.

As the atmosphere always contains carbon dioxide gas, particularly in a
laboratory where many gas flames are burning, considerable caution is neces-
sary in testing for a small quantity of carbonic acid. In the first place, if the
substance does not effervesce with acid, there is no use in making the test.
Moreover, carbon dioxide is odorless and it is absurd to think that effervescence
implies the presence of a carbonate when the escaping gas has the odor of hydro-
gen sulfide or sulfur dioxide.

A simple method of testing is as follows: Place about 1 em. of the powdered
solid in a test-tube, cover it with about 10 cc. of water and boil for about a
minute. This serves to expel the air from the substance and from the water.
Add a little 6-normal hydrochloric acid and watch closely to see if there is any sign
of effervescence. If there is effervescence notice whether the escaping gas has
any odor. . Dip two stirring rods into barium hydroxide solution, place one of the
rods between the second and third finger and the other between the third and
fourth and hold one of the rods inside the test-tube, without touching the sides,
and the other rod outside. If a carbonate is present, the barium hydroxide on
the rod inside the test-tube will become turbid faster than that on the rod held
in the air. This test is of no value, however, in the presence of a sulfite which
will also cause barium hydroxide to become turbid.

The above test naturally fails to detect traces of carbonate with certainty,
but it is satisfactory for all ordinary work except when a sulfite is present.
To detect traces of carbonate, even in the presence of a sulfite, the following
procedure is useful. Fit a 100 cc. flask with a rubber stopper containing two
holes. Through one hole insert a small dropping funnel so that it reaches nearly
to the bottom of the flask. Through the other hole insert a right-angled glass
tube which serves to lead the escaping gas to a second flask, containing about
5 gms. of chromic acid anhydride, CrOs;, dissolved in a little water and 25 ce. of
6-normal sulfuric acid. This will serve to oxidize any sulfurous acid or hydrogen
sulfide that may be set free. Arrange the tubes in this second flask so that the
gas passes down to the bottom of the solution and leaves the flask through
an exit tube which just reaches below the rubber stopper. Connect this flask .
in the same way with a third flask and connect the exit tube from this flask with
a drying tube filled with soda-lime, to prevent carbon dioxide getting in from
the air. Place the powdered substance in the first flask and cover it with
25 cc. of water. With the third flask empty, conduct a stream of air free from
carbon dioxide through the apparatus for ten minutes while heating the
water in the first flask. This is best accomplished by applying suction at
the end of the train and drawing the air through soda-lime in a drying tube
which is placed in a rubber stopper that fits the neck of the dropping funnel.
Instead of soda-lime, caustic potash solution (1 : 2) may be used to:remove the
carbon dioxide from the air. When the carbon dioxide has been expelled from
the apparatus, close the stop-cock in the dropping funnel, take away the flame.
and quickly add about 25 cc. of barium hydroxide solution to the last flask.
With the apparatus all connected, introduce 25 cc. of 6-normal sulfuric acid

CARBONIC ACID — 355

into the dropping funnel and allow it to run slowly into the flask containing the
substonce. If not enough gas is evolved to produce a turbidity in the barium
hydroxide solution, apply gentle suction as before, after a slight vacuum has
been produced, carefully open the stop-cock of the dropping funnel, with the
soda-lime tube in place, and gradually heat the sulfuric acid to boiling. Con-
tinue drawing air through the apparatus for fifteen minutes if necessary. In
case a very slight turbidity is obtained in the barium hydroxide solution it is
best to run a blank on the apparatus with all the reagents, and then repeat
the experiment. If the test does not give a decided result, the presence of
carbonate should never be reported.

2. Concentrated Sulfuric Acid reacts in the same way as dilute
sulfuric acid, only more violently.

3. Silver Nitrate precipitates white silver carbonate, which be-
comes yellow on the addition of an excess of the reagent. On boiling
with considerable water, the carbonate is partly decomposed into
brown silver oxide and carbon dioxide; but the carbonic acid is not
expelled completely except by heating to 200°. Silver carbonate is
very soluble in ammonia and in nitrie acid.

4. Barium Chloride precipitates white, voluminous barium car-
bonate, in the cold, which gradually on standing, but more quickly
on warming, becomes crystalline and denser.

Behavior of Carbonates on Ignition

The carbonate of the alkalies melt with but slight decomposi-
tion. Barium carbonate is not decomposed on charcoal before the
blowpipe, and does not melt; only at a white heat is it decomposed
into infusible barium oxide and carbon dioxide. All remaining
carbonates are decomposed at the temperature of the blowpipe into
oxide and carbon dioxide. The oxides of the noble metals are decom-
posed further into metal and oxygen.

PERCARBONIC ACID, HeC20¢

Free percarbonic acid is not known, but its potassium salt is stable in the
dry state. When exposed to moisture it decomposes into hydrogen peroxide
and potassium bicarbonate: |

K.C.0,+2H,0 = H,0O.+ 2KHCO:3.

If the salt, which is characterized by its pale blue color, is placed in con-
siderable cold, dilute sulfuric acid, it dissolves with evolution of carbon dioxide
and formation of potassium acid sulfate and hydrogen peroxide. The solution
gives all the characteristic reactions of hydrogen peroxide.

To distinguish between percarbonic acid and hydrogen peroxide, dissolve

10 gms. of potassium iodide in water and add 0.1 to 0.3 gm. of the finely
powdered substance. If potassium percarbonate is present, iodine is at once
liberated:

C.0.~+2I- —> 2CO;~+I2._

If only potassium bicarbonate and hydrogen peroxide are present, the liberation
of iodine will take place much more slowly.

Silver nitrate and barium chloride when treated with a percarbonate, give —

white precipitates which are soluble in dilute nitric acid.

eee

~~) oe A

BORIC (BORACIC) ACID, H3BO3

Occurrence.—Boric acid is found native as sassolite; in the form
of its sodium salt, as borax or tinkal, NazB407-10H20; as boracite,
2Mg3Bs015+MgCle; and in many silicates, such as axinite, tour-
maline, datolite, etc.

_ Crystallized borie acid forms colorless plates, with a mother-of-
pearl luster, which are soluble in water (100 parts water dissolve
4 parts of boric acid at 15°, and 33 parts at 100°). The aqueous
solution reacts acid, and is a poor conductor of electricity.

By heating boric acid to 100°, it loses one molecule of water and
is changed to metaboric acid, HBOz. The latter loses more water
when heated to 160°, forming pyroboric acid, H2Bs407; which, on
ignition, loses all its water, being changed to the anhydride of boric
acid, boron trioxide, which remains as a difficultly-volatile, hygro-
scopic glass.

The salts of boric acid, the borates, are derived from the meta-
and pyroboric acids. The salts of the ortho acid, H3BOs3, are not
known in the pure state.

In a few exceptional cases boron acts as a metal, forming B(HSO;);, (BO).SO,

BF;,BPOx,, etc. The last compound is insoluble in water and dilute acids, but
dissolves readily in caustic alkalies.

Solubility of Borates—The borates of the alkalies dissolve in water,
and the solution reacts alkaline.

A concentrated solution of borax behaves as if it contained sodium
metaborate, free boric acid, and a small amount of caustic alkali:

Na2eB407+3H20 @ 2NaBO2+2H3BO3;
NaBOo+ 2HOH @ NaOH+H3BOs.

The more dilute the solution, the greater the extent to which the
hydrolysis represented by the second equation will take place; so
that a very dilute solution of borax will react as if it contained simply
sodium hydroxide and free boric acid.

A solution of an alkali borate will behave differently towards
reagents, therefore, according to its concentration and temperature.
The remaining borates are difficultly soluble in water, but readily
soluble in acids and in ammonium chloride solution.

357

358 REACTIONS OF THE ACID CONSTITUENTS

REACTIONS IN THE WET WAY

For these reactions use a borax solution.

1. Dilute Sulfuric Acid.—No reaction.

2. Concentrated Sulfuric Acid.—No visible reaction. Most bo-
rates are decomposed by sulfuric acid, setting free boric acid, and
the latter is capable of coloring the non-luminous gas-flame with
a characteristic green tinge.

If, therefore, a little solid borate is placed in the loop of a platinum wire,
moistened with concentrated sulfuric acid, and heated at the edge of the
Bunsen flame, the characteristic green coloration will be noticed.

A great many natural silicates containing boric acid, when tested in the
above manner, will not give this flame coloration. To produce this coloration
moisten the mineral with hydrofluoric acid, place a little of it in the loop of a

platinum wire, and heat at the outer edge of the flame; if a borate is present the
latter will be colored distinctly green, owing to the formation of volatile boron

fluoride.
The presence of copper or of barium interferes with this test.

3. Concentrated Sulfuric Acid and Alcohol.—If an alkali or alka-
line-earth borate is treated in a platinum crucible with methyl alcohol,
then with concentrated sulfuric acid, the mixture stirred and the
alcohol lighted, a green-bordered flame will appear, due to the forma-
tion of boric acid methyl ester, B(OCHs)3.

4, Silver Nitrate produces, in moderately-concentrated, cold borax
solutions, a white precipitate of silver metaborate:

NazgB407 + 3HeO + 2AgNOsz = 2NaNO3 ~1. 2H3BO3 +2AgBOo.
On warming, a brown precipitate of silver oxide is obtained:
2AgBO> +3H.O = 2H3BO3 +Ag20.

From very dilute solutions, in the cold, silver nitrate produces a
brown precipitate oi silver oxide.

Silver borate is soluble in ammonia and in nitric acid.

5. Barium Chloride produces, in fairly concentrated solutions, a
white precipitate of barium metaborate:

NaeBs407 +BaClo +3H:2O a 2NaCl +- 2H3BO3 +Ba(BOzg)2,

soluble in an excess of barium chloride and in ammonium chloride. _

6. Calcium and Lead Salts behave similarly to barium chloride.

7. Turmeric.—If a piece of turmeric paper is placed in a solution
of free boric acid, apparently no change will take place unless consider-
able quantities of boric and sulfuric acids are present, but if the
paper is dried, it becomes reddish brown. If the brown paper is

BORIC ACID 359

again dipped in the solution of boric acid, the color remains; which
is also true if the paper is dipped in a dilute sulfuric or hydrochione
acid solution (difference from the alkali test with turmeric paper)
If the reddish-brown paper is moistened with caustic soda or potash
solution, the paper becomes bluish black; or, if only a small amount
of boric acid is present, grayish blue.

The shade and intensity of the color varies both with the amount of tur-
meric and with the amount of boric acid; with a very little boric acid, turmeric,
and very dilute caustic soda solution a nearly pure violet color is obtained and
with considerable boric acid and stronger alkali, a greenish-black color.

Borate solutions, when acidified with dilute hydrochloric acid, give the above
reaction. This sensitive and convenient test for boric acid must be used with
caution, for acid solutions of zirconic, titanic, tantalic, niobic, and molybdie
acids also color turmeric paper brown.

The reaction is much more sensitive if, instead of using the turmeric paper
itself, an alcoholic solution of turmeric is used. Place 2 or 3 drops of the yellow
solution in a porcelain dish, add the solution to be tested for boric acid, acidify
with acetic acid, and evaporate.to dryness on the water-bath. If-as much as
0.02 mgm. of BO; is present, the residue is colored a distinct reddish brown,
while 0.002 mgm. suffices to cause a visible reaction. (IF. Henz.)

8. Mercuric Chloride produces a red precipitate of basic mercuric
salt. If considerable free boric acid is present there is no precip-
itation.

Behavior of Borates on Ignition .

The hydrated borates of the alkalies melt with effervescence,

forming a colorless glass.

This glass has the property of dissolving many metallic oxides
when heated, whereby the often very characteristically colored meta-
borates are formed (borax beads); thus copper oxide is dissolved,

forming a blue glass:
NazBsO7+ Cu0 = 2NaBO2+Cu(BO2)s.

If this bead is heated in the reducing flame (i.e., with carbon)

two things can happen:
(a) The colored cupric salt is reduced to colorless cuprous salt:

4NaBO2+2Cu(BO2)2 +C= CO+Na2B407+2NaB02+Cuz(BO2)2.

(b) The cupric salt is reduced to. metallic copper, so that the bead
appears reddish brown and opaque:

4NaBOo+2Cu(BO2)2+C =CO2 T +2Na2Bs074+ 2Cu.

OXALIC ACID, H2C204

Occurrence and Preparation.—Oxalic acid occurs, in the form of
its acid potassium and calcium salts, in the sap of many plants.

It is prepared in large amounts by fusing sawdust with caustic
alkali. The resulting potassium salt is precipitated with milk of
lime, forming the insoluble calcium salt; and the latter is decom-
posed with sulfuric acid. Oxalic acid is also formed by the oxida-
tion of innumerable organic substances (such as sugar, starch, cellu-
lose (paper), by means of concentrated nitric acid.

It crystallizes from aqueous solutions in the form of colorless mono-
clinic prisms, H2C204-2H20. ;

By allowing the hydrated acid to stand over sulfuric acid the
water is lost, and the anhydrous acid remains, which, when heated —
to about 150° C., sublimes, forming needles. If heated still higher
it is completely decomposed into water, carbon dioxide, and carbon
monoxide:

HeC204= H20+COez FT +CO 7.

The crystallized, hydrated acid is soluble in water, alcohol, or
ether: 100 parts water at 20° dissolve 11.1 parts oxalic acid; 100
parts alcohol at 15° dissolve 33.2 parts oxalic acid; 100 parts ether
at 15° dissolve 1.5 parts oxalic acid.

Oxalic acid is a fairly strong, dibasic acid, and forms neutral
and acid salts, e.g., potassium oxalate, K2C204; potassium bin-
oxalate, KHC204; potassium tetroxalate, KHC204, HeC204:2H20.

Solubility —The oxalates are mostly insoluble in water, with the
exception of the oxalates of the alkalies and of magnesium. In an-
excess of an alkali oxalate many of the insoluble oxalates dissolve.
Thus ink spots and rust spots can often be removed from clothing
by means of a solution of oxalic acid or of potassium tetroxalate:

Fe203+6H2C204 > 2Ha[Fe(C204)]s +3H20.

Oxalice acid is also a good solvent for ferric phosphate. All oxalates
dissolve readily in mineral acids.

REACTIONS IN THE WET WAY

A solution of ammonium oxalate may be used for the following
reactions:

360

OXALIC ACID 361

1. Dilute Sulfuric Acid.—No reaction.

2. Concentrated Sulfuric Acid, on warming, acts as a dehydrating
agent, causing the evolution of equal volumes of carbon monoxide
and carbon dioxide; the latter burns with a blue flame:

H2C204 > H20+CO T +COe f..

In the presence of manganese dioxide, all oxalates evolve CO, with dilute
H.SO,: ;
H,C,0,+MnO,+H,SO, =MnSO,+2H,0+2CO, fT.

In the same way CO, is given off by the action of KMnQ, and dilute H,“O,
at about 60° C.:

2KMn0,+5H.C,0,+3H,80, =K,80,+2MnSO,+8H,0+10C0, ft.
3. Silver Nitrate precipitates white, curdy silver oxalate,
C204 +2Ag” =AgeC20z,

almost insoluble in water, but readily soluble in ammonia and in

- nitric acid.

4. Barium Chloride precipitates white barium oxalate soluble in
oxalic and acetic acids.

5. Calcium Chloride precipitates white calcium oxalate, insoluble
in oxalic acid, ammonium oxalate, and acetic acid, but readily soluble
in hydrochloric and nitric acids. It is the most insoluble of all oxalates.

6. Lead Salts precipitate white lead oxalate, soluble in nitric acid.

Behavior of Oxalates on Ignition

~All oxalates are decomposed on ignition with slight carboniza-
tion. The oxalates of the alkalies and alkaline earths are changed
to carbonates, with evolution of carbon monoxide. Stronger ignition
causes the formation of more or less oxide, the alkaline earth carbonate
being quantitatively changed to oxide by heating over the blast
lamp. The oxalates of the noble metals, and of iron, nickel, cobalt,
copper, etc., leave the metal itself; the oxide is formed in the case of
other metals.

TARTARIC ACID, H2C4H.10,

Occurrence.—Tartaric acid occurs partly free and partly as its acid
potassium salt in many fruit saps, particularly in that of the grape.

The free acid crystallizes in clear, monoclinic prisms, without
water of crystallization. Its aqueous solution is optically active,
turning the plane of polarized light to the right.

Three other modifications of this acid exist, possessing the same chemical
formula, but differing in their physical properties. One of these turns the plane
of polarized light to the left, and the other two are optically inactive.

Tartaric acid is very readily soluble in water (100 parts water
dissolve 132 parts of tartaric acid at 15°) and alcohol, but it is insol-
uble in ether. The salts are called tartrates.

Solubstlity—The neutral alkali tartrates are very soluble in water,
as also is acid sodium tartrate, while the acid potassium and the acid
ammonium tartrates are difficultly soluble in water.

The remaining tartrates are difficultly soluble in water, but all
dissolve, more or less readily, in neutral alkali tartrate solution, form-
ing complex salts.

The most important commercial salts of this acid are ‘‘ cream
of tartar,’ KHC4H40¢, “ Rochelle salt,’ KNaC4H40¢, and “ tartar
emetic,’ K(SbO)C4H40¢.

REACTIONS IN THE WET WAY

A solution of Rochelle salt (sodium potassium tartrate) may be
used for these reactions.

1. Dilute Sulfuric Acid.—No reaction.

2. Concentrated Sulfuric Acid causes carbonization on warming,
with evolution of sulfur dioxide.

3. Silver Nitrate produces no precipitation in a solution of free
tartaric acid, but in the solution of a neutral tartrate, a white, curdy
precipitate is immediately formed,

C4H40¢6 + 2Agt — AgeC4H10.,,

readily soluble in nitric acid and in ammonia. By warming the
ammoniacal silver solution, metallic silver is deposited. This very
362

TARTARIC ACID } 363

important reaction for the detection of tartaric acid is performed
in the following manner:

Treat the pure tartrate solution with silver nitrate solution until no further
precipitation takes place, then add dilute ammonia drop by drop until the pre-
cipitate just dissolves. Place the test- tube containing the solution in water
which has been heated to 60-70° C. After standing for about fifteen minutes,
the silver will be deposited in the form of a beautiful mirror on the sides of the
test-tube. This very delicate reaction cannot be performed with certainty in
the presence of other acids. In this case the tartaric acid should first be
' precipitated as potassium acid tartrate. Concentrate the solution to a small
volume, add a little solid potassium carbonate, acidify with strong acetic
acid and stir the cold solution vigorously; a precipitate of potassium acid
tartrate will form at once if considerable tartrate is present. Filter off the
precipitate, wash with a little cold water, and dissolve it in as little caustic
soda solution as possible. In this way a solution is obtained which will
readily give the silver mirror on the addition of silver nitrate and treatment as
above.

If no precipitate is formed -on the addition of the acetic acid, add a little
alcohol, which causes the precipitate to form more readily. Filter off the
precipitate, wash with diluted alcohol, dry, dissolve in dilute sodium hydrox-
ide, and treat as above. If the alcohol is not removed by drying, a mirror is
sometimes formed when tartaric acid is absent.

4. Calcium and Barium Chloride.—If to a concentrated solution
of neutral alkali tartrate, in the absence of ammonium salts, calcium
chloride solution is added drop by drop, a white amorphous pre-
cipitate is formed which redissolves, forming soluble calcium tartrate
anions:

2C4H4106 +Ca — [Ca(C4H40¢)o]-.

Only after the addition of enough calcium chloride to decompose
completely the alkali tartrate is a permanent precipitate formed,
which at first is flocculent, but soon becomes crystalline, consisting
of neutral calcium tartrate:

[(Ca(C4H40¢)2]> +Cat* — 2CaCsH4O¢.

In dilute solutions the first addition of calcium chloride often produces no
precipitation; but after standing some time (or more quickly on rubbing the
sides of the test-tube with a glass rod) the crystalline precipitate is deposited,
CaC.H,O.+4H.0. Calcium tartrate is very difficultly soluble in water; 100
parts water at 15° C. dissolve 0.0159 part of the crystalline salt, and 100 parts
boiling water dissolve 0.0285 part of the salt. The precipitate is soluble in
acetic acid (difference from calcium oxalate) and also in a solution of concen-
trated caustic alkali (free from carbonate), probably forming a complex ion:

2CaC,H,0O,+20H- — H,0+[(CaCyH.Q¢),0]-.

364 REACTIONS OF THE ACID CONSTITUENTS

On boiling this solution, calcium tartrate is reprecipitated in the form of a
voluminous gelatinous precipitate, which again goes into solution on cooling..
The presence of ammonium chloride retards the formation of the calcium tar-
trate, but does not prevent it; after standing some time, the precipitate settles
out in the form of a heavy crystalline powder (difference from citric acid).

5. Potassium Salts produce no precipitation in neutral solutions
of alkali tartrates; but if the concentrated solution is acidified with
acetic acid, a precipitate of crystalline potassium acid tartrate is
formed upon stirring vigorously:

C4H40¢6-+Kt+Ht — KHC4H10¢.

Potassium acid tartrate is difficultly soluble in water (100 parts water
dissolve 0.45 part of salt) and in acetic acid, but is readily soluble in mineral
acids or in caustic alkali and alkali carbonate solutions. The precipitate may be
dissolved in a little ammonia and the above test with calcium chloride obtained.

If a concentrated solution of free tartaric acid is treated with potassium
chloride, a precipitate of potassium acid tartrate is formed in spite of the pres-
ence of the hydrochloric acid which is set free. From dilute solutions the pre-
cipitate appears only after adding sodium acetate (cf. p. 79). :

The presence of considerable boric acid greatly interferes with the formation
of the potassium acid tartrate. In such cases add ammonium chloride to the
concentrated solution in a test-tube, then some calcium chloride solution, and
rub the sides of the glass. with a stirring rod. If only a little tartaric acid is
present it may be necessary to let the solution stand twenty-four hours. Filter
off any precipitate that may form, wash it two or three times with alcohol, and -
dissolve the calcium tartrate in a little freshly-prepared potassium hydroxide
solution (1:5). Filter and heat the filtrate to boiling. Amorphous calcium
tartrate should be precipitated. Another way to identify the tartrate in the
first precipitate of calcium tartrate is as follows: Place the washed precipitate
in a test-tube, add a crystal of silver nitrate, a few drops of 6-normal ammo-
nium hydroxide, and place the test-tube in hot water; a distinct silver mirror
should form if a tartrate is present.

6. Lead Acetate produces in neutral solutions a white, flocculent
precipitate of lead tartrate, easily soluble in nitric acid and in ammonia.

7. Magnesia Mixture. If a concentrated tartaric acid solution is
treated with an excess of magnesia mixture, 10 ce. of strong ammonia
and a volume of alcohol equal to that of the solution, then, after
shaking and allowing to stand twelve hours, the tartaric acid is pre-
cipitated quantitatively as crystalline, basic magnesium tartrate
insoluble in 50 per cent alcohol (difference from malic and succinic
acids) :

C4H106—>+2Mg* *+20H~+H20 — Mgo(OH)2 (C4H40¢)- H20.

Filter off the precipitate obtained in the above test, wash it with 50 per
cent alcohol, dry and transfer the precipitate to a test-tube with the aid of a

' TARTARIC ACID 365

glass rod. Add a little silver nitrate solution, a slight excess of 6-normal
ammonium hydroxide and heat to about 60°. A mirror should form if a tar-
trate is present.

REACTIONS IN THE DRY WAY

If tartaric acid is heated to 135° C., it melts, and on stronger
ignition it is decomposed, leaving a residue of carbon and giving off
empyreumatic odors (smell of burnt sugar).

The alkali tartrates are also decomposed by ignition, leaving a
residue of carbon and alkali carbonate, which effervesces on treat-
ment with acid. |

Ammonium tartrate leaves a residue of carbon, which does not
effervesce on treatment with acids. The tartrates of the alkaline
earths leave behind a mixture of carbon and carbonate; on very
strong ignition the latter is changed to oxide.

The tartrates of those metals whose oxides are reduced by car-
bon are left in the form of metal (Ag, Pb, Fe, Ni, Co, ete.).

CITRIC ACID, H3C 6Hs5O07

Citric acid is found in nature in the juices of many fruits. It is
a tribasic acid, readily soluble in water and in alcohol, but difficultly
soluble in ether. Its salts are called citrates.

Solubility —The citrates of the alkalies are soluble in water, and
form, with the insoluble citrates of the heavy metals, very soluble
complex salts, whose solutions are not precipitated by alkali hydrox-
ides, alkali «carbonates, ammonia, etc.

REACTIONS IN THE WET WAY

A solution of potassium citrate may be used.

1. Dilute Sulfuric Acid.—No reaction.

2. Concentrated Sulfuric Acid on being heated with a citrate, as
with most a-hydroxyacids, causes formic acid, HCO2H, to be formed,
which then’ breaks down into water and CO; at the same time some
acetone dicarbonic acid, (CH2)2CO-(COe2H)s2, is formed and this breaks
down into acetone, (CH3)2CO, and COs.

H3C.Hs07 =? HCO2H+(CH2)2CO(CO2H )2,
HCO2H — H20+C0 T, (CH2)2CO’COsH)2 —.(CH3)2CO+2C02 T .

A part of the citric acid is carbonized and this causes reduction of the -
sulfuric acid so that some, SQOz2 is evolved.

3. Silver Nitrate produces in neutral solutions a flocculent pre-
cipitate of silver citrate, AgsCe6HsO7, readily soluble in nitric acid
and in ammonia. On heating the ammoniacal solution to 60° C.,
no silver mirror is formed; but on heating the solution to boiling,
the silver is gradually deposited.

4. Barium and Calcium Chloride give no precipitation in neutral
solutions (difference from tartaric acid). If, however, caustic soda
solution is added to the solution which contains an excess of calcium
chloride, a flocculent precipitate of tertiary calcium citrate is at once
formed, insoluble in caustic alkali, but readily soluble in ammonium
chloride. On boiling the solution in ammonium chloride, crystalline
calcium citrate is precipitated, which is now insoluble in ammonium
chloride.

366

CITRIC ACID 367

5. Lime Water in excess produces no precipitation in cold solutions
of neutral citrates; on boiling, there is formed a flocculent precipitate
of calcium citrate, which almost entirely redissolves on cooling.

_ 6. Lead Acetate precipitates from solutions of the free acid, and
those containing neutral salts, amorphous Pb3(CsH;07)2-H2O.

7. L. Stahre’s Test for Citric Acid.* To the solution of free citric acid in
water, or to the solution of a citrate in very dilute sulfuric or nitric acid (not
hydrochloric) add 2 to 5 drops of tenth-normal permanganate solution and heat
a short time at 30° to 40° (the solution must not boil!). As soon as the solution
is colored brown, or becomes turbid by the precipitation of a little manganese
dioxide, add 1 or 2 drops of ammonium oxalate solution and about 1 ce. of
10 per cent sulfuric acid, which will clear up the solution. Now, add a few
drops of bromine water, and a distinct, crystalline precipitate of pentabromace-
tone will be obtained. The bromine water may also be added before the per-
manganate solution and sometimes the results are better.

This test is so sensitive 0.38 mgm. of citric acid in 1 ec. of water can be
detected.

The experiment succeeds in the presence of tartaric, malic, oxalic, sulfuric
and phosphoric acids, except that a little more permanganate is required.

In the Stahre test the following reactions take place:

(a) The permanganate oxidizes the citric acid to acetonedicarboxylic acid
with evolution of carbon dioxide:

€H;CsH;0;+2Mn0, +€Ht — 5CO, T +2Mn*t++£H.0+5(CH,)2,CO(CO.H)..

(6) The acetonedicarboxylic acid reacts with bromine, forming penta-
bromacetone:

(CH,),CO(CO.H).+5Br, a 2CO, T +5HBr+C.HBr,CoO.

If the permanganate is allowed to act longer upon the citric acid the acetone.
dicarboxylic acid is converted gradually into acetone, the reaction taking
place more quickly on boiling:

(CH,).CO(CO.H)2 > 2CO, T +(CH:;).CO.

Acetone itself is not brominated as readily as the acetonedicarboxylie acid,
and for this reason care should be taken not to let the temperature rise above
40° during the treatment of the citric acid with permanganate.

The citrates on treatment with bromine, without previous oxidation with
permanganate, will also give pentabromacetone:

K;C.H;0; + 6Brz =3KBr + 3CO, T + 4HBr+ C.HBr;CO,

8, Mercuric Sulfate.—Denigés’ reagent.t Dissolve 5 gms. HgO in 100 ee.
of water and 20 ee. cone. H.SO,. Treat the solution of the citrate with 1/20
as much reagent and heat to boiling, then add a few drops of 0.1N KMnO,
solution. A white crystalline precipitate is formed.

*L. Sraure, Z. anal. Chem., 36 (1897), 195; also ALFRED WOuLK, ibid., 41,
94 (1902).
{ Comptes rend., 188, 32; Z. anal. Chem., 88, 718 (1899); and 40, 121 (1901).

368 REACTIONS OF THE ACID CONSTITUENTS

The precipitate has the composition: Hg;0.80,-2[(CH2),CO(CO:).JHg and
is a mixture of basic mercuric sulfate and the mercuric salt of acetone dicar-
boxylic acid. The reaction is very sensitive and enables one to detect 0.5 gm.
of citric acid dissolved in a liter of water. The reaction is, however, not pecu-
liar to citric acid but is shown by many other ketonic compounds, |

REACTIONS IN THE DRY WAY

The citrates, on ignition, behave exactly like the tartrates

PHOSPHOROUS ACID, H3PO3

-Formation—By the slow combustion of phosphorus in the air
phosphorus trioxide is formed which, as it is the anhydride of phos-
phorous acid, reacts with cold water to form the acid:

P203+3H20 =2H3PO3.

Phosphorous acid is formed much more readily by the action of
water on the trihalogen compounds of phosphorus:

PCl3+3HOH =3HCl-+ H3P0Os.

The hydrochloric acid is removed by evaporation, and the last
traces of uncombined water by heating to 180°. If the mass is then
allowed to cool, it solidifies to a crystalline, hygroscopic substance
which melts at 70°.

By neutralizing the solution of phosphorous acid with bases; the
phosphites are obtained. It is never possible, however, to replace
more than two of the hydrogen atoms with metal; so that phosphorous
acid is considered a dibasic acid. Certain organic compounds are
known, however,-which are derived from tribasic phosphorous acid,
H3POs3.

Solubility Only the phosphites of the alkalies are soluble in
water, but they are all soluble in acid.

REACTIONS IN THE WET WAY

A solution of sodium phosphite should be used.

1. Dilute Sulfuric Acid.—No reaction.

2. Concentrated Sulfuric Acid causes no reaction in the cold; on
heating, the phosphorous acid reduces the sulfuric acid to sulfurous
acid, easily recognized by the odor of burning sulfur,

H3PO3-+H2S04 = H3P01+H20+80z2 f .

3. Silver Nitrate produces at first a white precipitate of silver
phosphite,

HPO3-+2Ag* — AgoHPOs,
369

370 REACTIONS OF THE ACID CONSTITUENTS

which in the case of a concentrated solution is changed in the cold
to metallic silver; while in dilute solutions this reduction takes place
only on warming:

AgoHPO3; -}- H2O = H3P04 +2Ag.

4. Barium Chloride precipitates white barium phosphite, soluble
in all acids.

5. Lead Acetate precipitates white lead phosphite, insoluble in
acetic acid.

6. Mercuric Chloride is te reduced phosphorus acid in
the cold, but more quickly on warming, to mercurous chloride:

2H¢gClo+H3P03+H20 = H3P01+2HCl+Hg2Ch.

If the phosphorous acid is present in excess, the reduction in the
hot solution (not in the cold) goes further, and gray metallic mercury
is deposited:

HgeCle a. H3PO3 +H.O = HsPO4 -f- 2HCl +2H¢g.

7. Nascent Hydrogen (zinc and sulfuric acid) reduces phosphor-
ous acid to phosphine:

H3P03+6H =3H20+PHs} .

If the phosphine is allowed to act upon a concentrated solution
of silver nitrate (1:1), or better still, upon the solid silver nitrate,
the latter is colored yellow, as with arsine:

PH3+6AgNO3 = PAgs3-3AgNO3+3HNOs.

By the addition of water this yellow compound is decomposed
with separation of grayish-white silver:

PAgs . 3AgNOz3 +3H2O = 3HNOs: +H3P03 +6Ag.

The phosphorous acid is, however, immediately oxidized by the
nitric acid to phosphoric acid:

3H3P03+2HNO3 = He0+2N0+3H3POs4.
The mixture of phosphine and hydrogen burns with an emerald-

green flame.

' 8. Sulfurous Acid is reduced by phosphorous acid to hydrogen
sulfide:

3H3PO03+ H2S03 =3H3P04+ HS f .

PHOSPHOROUS ACID 371
9. Concentrated Potassium Hydroxide Solution changes a phos-
phite to phosphate, with evolution of hydrogen,
KeHPO3+ KOH = K3P04+ He fF ,
but with dilute caustic potash the hydrogen evolution is very slight.

REACTIONS IN THE DRY WAY

By ignition, phosphorous acid (like hypochlorous acid) is changed
at the cost of its own oxygen to the higher compound, while the oxidiz-
ing part of the acid is reduced to its hydrogen compound:

3HCI1O = HClO3+2HCl;
4H3PO3=3H3P04+PHs f .
The phosphates behave similarly:
SNazHPO3 =4NazgP04+ NasP207+H20+2PHs 7 .

METAPHOSPHORIC ACID, HPO;

The monobasic metaphosphoric acid is obtained by treating phos-
phorous pentoxide with cold water,

P205;+H20 =2HPOs,
and also by the strong ignition of orthophosphoric acid:
H3P04= H20+HPOs.

Metaphosphoric acid is a colorless, glassy, hygroscopic mass.
On boiling its aqueous solution, or slowly in the cold, it adds water to-
the molecule, and is changed to orthophosphoric acid:

HPO3+ H20 = H3P04.

The metaphosphates are readily obtained by heating the mono-
metallic salts of orthophosphoriec acid,

NaH2PO4=Hz20 fT +NaPOs,
or by igniting sodium ammonium phosphate: |
NaNH4HPO4=H20 , +NH3 T +NaPOs.

The meta salts are changed into orthophosphates by boiling the
aqueous solution in the presence of mineral acid.

Solubility.—The metaphosphates of the alkalies and of magnesium
are soluble in water; the remaining salts are difficultly soluble or
insoluble in water, readily soluble in nitric acid, and in an excess of
metaphosphoric acid or an excess of alkali metaphosphate.

REACTIONS IN THE WET WAY

Sodium metaphosphate is used for the following tests:

1. Sulfuric Acid causes no visible reaction. .

2. Silver Nitrate precipitates white silver metaphosphate, soluble
in ammonia and in nitric acid:

PO3--+Agt — AgPOs.

3. Barium Chloride precipitates voluminous barium metaphos-
phate, soluble in an excess of sodium metaphosphate, from which
372

METAPHOSPHORIC ACID 373

solution ammonia causes no precipitation. Barium sodium dimeta-
phosphate (or a similar polymetaphosphate) is probably formed.

4. Magnesium Salts cause no precipitate from moderately dilute
solutions, even on boiling (difference from orthophosphoric acid).

5. Ammonium Molybdate produces no precipitate in the cold;
but, on boiling the acid solution, metaphosphoric acid is changed to
orthophosphorie acid and the characteristic precipitate of ammonium
phosphomolybdate is formed. .

6. Albumin Solution is coagulated by an aqueous solution of the
free acid (difference from pyro- and orthophosphoric acids), but not
by a solution of alkali metaphosphate, except on the addition of acetic
acid.

7. Nascent Hydrogen does not reduce metaphosphoric acid (differ-
ence from phosphorous acid). |

Behavior on Ignition

The alkali metaphosphates, on being fused, form a glassy mass,
which has the property of dissolving many metallic oxides, forming
orthophosphates with characteristic colors. (See Phosphoric Acid.)
By fusion with soda, orthophosphates are formed from metaphos-
phates.

PYROPHOSPHORIC ACID, H4P207

The tetrabasic pyrophosphoric acid is formed by heating ortho-
phosphoric acid to 213°. It is a soft, glassy mass, readily soluble in
water; and in solution it gradually adds water to the molecule. and is
changed to phosphoric acid, the change taking place quickly on
boiling the solution.

The salts of pyrophosphoric acid, the pyrophosphates, are obtained
by igniting the dimetallic phosphates:

2Na2HP O4 = He2O +NaaP207.

Solubility—The pyrophosphates of the alkalies are soluble in
water; the remaining pyrophosphates are difficultly soluble or insol-
uble in water, but are all soluble in acids, and some are soluble in an
excess of sodium pyrophosphate.

REACTIONS IN THE WET WAY

A solution of sodium pyrophosphate is used for these tests.

1. Sulfuric Acid.—No reaction.

2. Silver Nitrate gives a white, curdy precipitate, soluble in am-
monia and in nitric acid.

3. Barium Chloride causes a white, amorphous precipitate, coer
in acids.

4. Magnesium Chloride produces a white precipitate which is
soluble in an excess of the magnesium salt, as well as in an excess
of sodium pyrophosphate. By boiling this solution a_ precipitate
is formed, which does not disappear on cooling.

5. Ammonium.Molybdate produces no precipitation in the cold;
but, on warming, yellow ammonium phosphomolybdate is precipi-
tated.

6. Albumin is not coagulated by free pyrophosphoric acid (differ-
ence from metaphosphoric acid). ;

BEHAVIOR IN THE DRY WAY

All pyrophosphates on being fused with sodium carbonate are
changed to orthophosphates:

NasP207+Na2CO3=CO2 T +2NagPO04.
374

IopIc ACID, HIO;

Occurrence.—In sea-water and in Chili saltpetre as potassium
iodate.

Formation.—By oxidizing iodine with fuming nitric acid or by
the action of chlorine upon iodine suspended in water:

312+10HNO3=6HIO3+10NO T +2H20;
Iz+6H20+5Cle = 10HC1+2HI103.

The most important iodate, KIOs3, is obtained by the action of
iodine upon a slightly acid solution of potassium chlorate:

5KC10O3+3I12+3H20 = 5KIO3+ HIO3+5HCI.

Iodates are also formed by the action of iodine upon alkali hy-
droxide solutions:

3le+6KOH = 5KI+ KIO3+3H20.

In alkaline solutions iodides are oxidized to iodates by hypo-
chlorites and potassium permanganate.

Solubility —The -iodates of the alkalies are soluble in water, but
the remaining iodates are difficultly soluble or insoluble.

REACTIONS IN THE WET WAY

1. Sulfuric Acid.—Neither dilute nor concentrated sulfuric acid
decomposes iodic acid; but if reducing substances are present at the
same time (such as hydriodic acid, hydrogen sulfide, ferrous salts, :
etc.), the iodic acid is reduced, with separation of iodine:

I03~+5I-+6H* — 3H20+43le.

2. Silver Nitrate precipitates white, curdy silver iodate, AgIOs,
readily soluble in ammonia, but difficultly soluble in nitric acid.

3. Barium Chloride precipitates white barium iodate, difficultly
soluble in hot water (100 parts of boiling water dissolve 0.6 part of
the salt), and only slowly soluble in dilute hydrochloric or nitric acids.

4. Lead Acetate precipitates lead iodate, difficultly soluble in
water and only slightly soluble in nitric acid.

375

376 REACTIONS OF THE ACID CONSTITUENTS |

5. Reducing -Agents.
(a) Hydriodic acid reduces iodic acid, with separation of iodine:

H1I03+5HI=3H20+3le.

If the solution is concentrated, the iodine separates out as a brown
powder; dilute solutions are colored yellow. The iodine may be
absorbed, with a reddish-violet color, by shaking the aqueous solution
with chloroform or carbon disulfide.

(b) Sulfurous acid also causes separation of iodine, unless a large —
excess of sulfurous acid is added: .

2103~+5S037+2H*t > 5804 +H20+]e,
IO3_+3803° — 3804-4 T-.

The reduction takes place according to the last reaction, when
three molecules of sulfurous acid are present to one of iodate; sul-
furous acid reacts with free iodine (cf. p. 349).

REACTIONS IN THE DRY WAY

Heated on charcoal the iodates deflagrate, but not so strongly
as the chlorates; they are all decomposed on being heated, some
with and some without the separation of iodine. Thus all neutral
iodates of the alkalies are easily decomposed into iodide and oxygen,
while the biiodates set free iodine at the same time:

2KIO3 =2KI+302T ;
4/KIO3-HIO3] =4KI+1102 fT +2H20+42Io.

GROUP IV

Silver Nitrate produces in neutral solutions a colored precipitate,
soluble in nitric acid.

Barium Chloride also produces a precipitate which is soluble in
nitric acid.

PHOSPHORIC ACID, H3P04

Orthophosphoric acid is obtained by the oxidation of phosphorus
by means of nitric acid, or by boiling the meta- and pyro-phosphoric
acids with water. It is a tribasic acid, and forms salts in which either
one, two, or three of its hydrogen atoms are replaced by metals (cf.
p. 10), e.g., NaH2PO4, NazHPO«4 and NagPOs..

Solubility—The phosphates of the alkalies are soluble in water,
and so are the primary salts of the alkaline earths. The secondary
phosphates of the alkaline earths are very difficultly soluble, while
the corresponding tertiary phosphates (as well as all other phos-
' phates) are insoluble. All phosphates dissolve in acids (cf. p. 47).

REACTIONS IN THE WET WAY

Use a solution of disodium phosphate for these reactions.

1. Sulfuric Acid, dilute or concentrated, produces no visible change.

2. Silver Nitrate produces a yellow precipitate of silver phosphate
(difference from meta- and pyrophosphoric acids),

2HPOs" +3Ag*=H2POs. +AgsPOu,

readily soluble in nitric acid and in ammonia. The precipitate, there-
fore, can be formed only in neutral solution.
3. Barium Chloride precipitates white, amorphous barium phos-
phate:
HPO.z-+Bat * — BaHPOs..

_ In the presence of ammonia the less soluble tertiary salt is precipitated:
2HPO4-+3Bat +1LONH3 oo 2QNHat+ Baz (PO4)e.

The barium phosphates (as well as those of the other alkaline
earths) are easily dissolved by acids, even acetic acid (difference from
377

378 REACTIONS OF THE ACID CONSTITUENTS

aluminium and ferric phosphates). From these acid solutions, am-
monia reprecipitates the phosphate.

In pure water H.PO,~ is dissociated to about 0.1 per cent, but in the presence
of acetic acid to a much less extent. BaHPO, in contact with water furnishes
more HPO; ions than does H.PO, in the presence of acetic acid, therefore
BaHPQ, dissolves. Similarly Bas(PO,)2 dissolves in order to establish equilib-
rium between Ht and PO; ions. Aluminium phosphate is less soluble than
barium phosphate and therefore requires a stronger acid to dissolve it.

By adding ammonia the hydrogen ions are neutralized and PO; ions are
obtained in equilibrium with NH,* ions, and Ba;(PO,)2 is reprecipitated.

4. Magnesia Mixture (an aqueous solution of ammonium chloride,
ammonia, and magnesium chloride) precipitates from very dilute

solutions white, crystalline magnesium ammonium phosphate,
MgNH4P04+6H20,

HPO. +Mg*tt+NH3 — MgNHyPO,,

which is soluble in all acids, but practically insoluble in dilute 23
per cent ammonia. This is a very sensitive reaction (cf. p. 75).

5. Ferric Chloride.—If a solution of sodium phosphate is treated
with ferric chloride, a yellowish-white precipitate of ferric phosphate
is formed: |

HPO,-+Fet ++ = Ht+FePQ,.

Hydrogen ions are formed in this reaction and, as might be expected from
the mass-action principle, the precipitation of the phosphoric acid is not quan-
titative unless the greater part of the hydrogen ions are removed. This may
be accomplished by adding ammonium acetate, as the hydrogen ions must
form non-ionized acetic acid to be in equilibrium with the acetate ions (cf.

. 46).
2 HPO,-+(C.H;0.-+Fet tt — HC.H;0.+FeP O,.

Moreover, if the reaction takes place in a boiling, dilute solution, the excess
of the iron can be precipitated as basic ferric acetate. If the solution is filtered
hot, a filtrate is obtained which is free from iron and from phosphoric acid. If,
however, it cools, some of the iron goes back into solution and, as ferric phos-
phate is appreciably soluble in ferric acetate solution, some of the phosphoric
acid also goes into solution.

Since ferric phosphate, unlike the phosphates of the alkaline earths, is
insoluble in acetic acid, it is evident that phosphoric acid may be removed
from a solution of alkaline earth phosphate in acetic acid by adding ferric
chloride, an excess of soluble acetate and boiling.

To accomplish this, dissolve the phosphate in as little hydrochloric acid as
possible, add ammonium carbonate until a slight permanent precipitate is
obtained and dissolve the precipitate by adding one or two drops of 6-normal
hydrochloric acid. Add an excess of ammonium acetate and ferric chloride,
drop by drop, until the solution above the yellowish-white precipitate of ferric
phosphate is colored distinctly brown by colloidal ferric hydroxide. Dilute

PHOSPHORIC ACID 379

with considerable water, heat to boiling and filter while hot. To detect phos-
phoric acid in the precipitate dissolve it in nitric acid, evaporate the solution
to a small volume and treat with ammonium molybdate solution; a yellow,
crystalline precipitate of ammonium phosphomolybdate proves the presence
of phosphoric acid. Or, dissolve the iron precipitate in hydrochloric acid, add
2 gms.of tartaric acid to prevent the precipitation of iron, add ammonia in excess
and then some magnesium-ammonium chloride mixture. A white precipitate
of magnesium ammonium phosphate shows the presence of phosphoric acid.

6. Ammonium Molybdate, in large excess, precipitates from nitric
acid solutions in the cold on standing (more quickly on slightly warm-
ing) a yellow, crystalline precipitate of ammonium phosphomolybdate:

H3P04 + 12(NHa4) 2M004+21HNO3=
= (NH,)3PO4-12M003+21NHiNO3+12H20.

This reaction is analogous to the reaction with arsenic acid (cf. page 231),
except that the arsenic compound is formed quickly only at the boiling tem-
perature. The presence of ammonium nitrate greatly facilitates the formation
of this precipitate.

Ammonium phosphomolybdate is readily soluble in alkalies and in ammonia,

(NH,);PO,-12M00,+280H- — 3NH.*+HPO+12Mo00,-+1 1.0,

also in an excess of alkali phosphate solutions, forming conipoulids which
contain less molybdenum. It is, therefore, always necessary to prevent the
formation of such compounds by the addition of a large excess of ammonium
molybdate.

Detection of Phosphorus in Iron and Steel.—Phosphorus is present in iron
and steel as iron phosphide, but only to a slight extent (usually less than 0.1
per cent). To detect the phosphorus it is necessary to oxidize it to phosphoric
acid and then use one of the above reactions. As, however, very small amounts
of phosphorus are present, it is necessary to start with a large amount of the
original substance in order to obtain a perceptible phosphorus test. It is best to
proceed as follows: Dissolve 5 to 10 gms. of the iron or steel in 60 cc. of 6-normal
nitric acid,* evaporate the solution to dryness and ignite over a free flame (with
constant stirring) until no more red fumes are given off. All organic matter is
thereby destroyed, silicic acid is dehydrated and the oxidation of the phosphorus
to phosphoric acid is completed. After cooling, dissolve the oxides in 50 ec. of
12-normal hydrochloric acid (warming gently), evaporate off the excess of acid,
dilute and filter off the silica. In the filtrate all the iron and all the phosphoric
acid will be found, and the latter may be detected by either the molybdate or
the magnesia-mixture reaction. To detect the phosphoric acid according to the
former method, evaporate to dryness the filtrate obtained after the removal
of the silica, dissolve the residue in as little 6-normal nitric acid as possible,
add 50 ec. of ammonium molybdate solution and 15-20 cc. of a 75 per cent

* If the iron were dissolved in HCl or H2SO,, part or even all of the phosphorus
would escape as phosphine. Nitric acid oxidizes nearly all of the phosphorus to
phosphoric acid.

380 REACTIONS OF THE ACID CONSTITUENTS

ammonium nitrate solution, heat the mixture gently, shake or stir vigorously
and allow it to stand an hour. A yellow, crystalline precipitate shows the
presence of phosphorus.

To detect the phosphorus according to the magnesia-mixture method, it is
necessary first to remove the greater part of the iron. Neutralize the hydro-—
chloric acid filtrate with ammonia, add a saturated solution of sulfur dioxide
and boil the solution, whereby the previously dark-colored solution is either
decolorized or becomes a light green. Add 20 cc. of 12-normal hydrochloric
acid, and boil the solution until the excess of sulfur dioxide is expelled. By
this operation all the ferric salt is reduced to ferrous salt. Add a few drops of
chlorine water (which forms a little ferric salt), neutralize with ammonia and
dilute to about a liter; add 3 cc. of a saturated solution of ammonium acetate,
5 ec. of acetic acid, and heat the solution to boiling. All the ferric salt and all
the phosphoric acid will be precipitated in the form of ferric phosphate and
basic ferric acetate, while the greater part of the iron remains in solution as
ferrous salt. Filter off the light brown precipitate through a small plaited
filter, wash it with hot water, and dissolve in dilute hydrochloric acid. Evapo-
rate the solution almost to dryness, add 2 gms. of citric (or tartaric) acid (which
should be dissolved in as little water as possible), add an excess of ammonia, and
precipitate the phosphoric acid by the addition of magnesia mixture. A white,
crystalline precipitate shows the presence of phosphoric acid.

7. Lead Acetate precipitates white lead phosphate, nearly insol-
uble in acetic acid:

2HPO4 + 3Pb**+ 2C2H302° — 2HC2H302+Pb3(PO4)o.

8. Nascent Hydrogen does not reduce phosphoric acid (difference
from phosphorous and hypophosphorous acids).

9. Metastannic Acid.—If metallic tin is added to a nitric acid
solution of phosphoric acid, or a phosphate, the tin is changed to
metastannic acid, which unites with the phosphoric acid, forming
an insoluble compound (probably a complex phospho-stannic acid).
This reaction is often used to separate phosphoric acid from other _
metals.

10. Mercurous Nitrate precipitates from solutions which are
almost neutral, white mercurous phosphate, soluble in nitric acid
but insoluble in acetic acid.

REACTIONS IN THE DRY WAY

The tertiary salts of the alkalies melt without decomposition;
the secoridary salts lose water and are changed to pyrophosphates
while, the primary salts form a glassy metaphosphate.

The so-called “salt of phosphorus,” or “ microcosmic salt,”
NaNH,HPO.+4H20, which is much used as a reagent, loses water

PHOSPHORIC ACID 381

and ammonia on being fused, forming a clear glass of sodium meta-
phosphate:

NaNH.HPO,-4H20 =5H20 ¢ +NHs f +NaPOs.

If the salt is heated in the loop of a platinum wire, a clear bead
is obtained—the so-called “‘ salt of phosphorus ”’ bead.

Just as metaphosphoric acid unites with water, on boiling its
solution, forming orthophosphoric acid,

HP O3 -}- H2O = H3POsz,

so sodium metaphosphate dissolves, at the fusion temperature, a
great many metallic oxides, forming characteristically colored ortho-
phosphates,

| NaPO3+Cu0 = NaCuP0Og (blue bead),
which may be changed in the reducing flame to metaphosphate again:
NaCuPO4+C=CO JT +Cu+NaP0Os.

Brownish red opaque bead

Many anhydrous phosphates are reduced by heating with mag-
nesium to phosphides, which, on being breathed upon, give the
peculiar odor of phosphine:

Cas3(POs)2+8Mg =8Mg0+CasP2;
CazP2+6H20 =3Ca(OH)2+2PHs f .

PHOSPHORUS, P. At. Wt. 31.04. Mol. Wt. Py =124.16. M.P.=44.1°

Occurrence.— Phosphorus is found in nature only in the form of
phosphates, of which calcium phosphate is the most important. It
occurs as apatite, Cas(PO4)3(Cl,F), in hexagonal crystals, and in an
impure state as phosphorite, which is used extensively as a fertilizer.
Calcium phosphate is also an important constituent of bones and the
seeds of plants.

A very interesting occurrence of phosphorus is pyromorphite
(ef. p. 205), isomorphous with apatite, vanadinite, and mimetesite.

Properties—Phosphorus exists in four allotropic forms: (a) As
ordinary or colorless phosphorus. (b) As red, crystalline phosphorus.
(c) As bright red phosphorus. (d) As black, crystalline phosphorus.

Ordinary phosphorus is poisonous, is colorless when pure (it be-
comes yellow on exposure to the light, and is coated with a layer of
red phosphorus), melts at 44° C., and ignites at 60° C. in the air,
so that it must be kept covered with water, in which it is insoluble.
It is readily soluble in carbon disulfide, and slightly soluble in ether.
It is easily oxidized by nitric acid to phosphoric acid:

3P4+20HNO3+8H20 = 12H3P04+20NO f.

The colorless phosphorus, but not the red modification, is oxidized
to hypophosphorous and phosphorous acids by exposure to moist
air. This causes the characteristic phosphorous odor, and, in the
dark, a pale green luminescence. If phosphorus vapors, or phos-
phine, are allowed to act upon most silver nitrate paper, the latter
‘is blackened, on account of the formation of silver phosphide and
metallic silver. The reaction probably takes place in this way: First,
the phosphorus reacts with water to form phosphine and hypophos-
phorous acid,

P4+6H20 =3H3P02+PHs f ,*
which then react with the silver nitrate:
H3P02+2H20+4AgNO3 =4HNO3+H3P04+4Ag;
PH3+3AgNO3 =3HNO3+AgsP.

* Phosphorus and water by themselves do not react in accordance with this
reaction, but it seems probable that they do in the presence of silver nitrate.
382

PHOSPHORUS 383

This exceedingly sensitive reaction for colorless phosphorus was
discovered by Scheuer.* It is a decisive test only when no other
substance is present, such as HeS, H3As, HgSb, formaldehyde or
formic acid, which is capable of reacting with silver nitrate.

Red phosphorus is crystalline (hexagonal, rhombohedral), and
is formed by heating ordinary phosphorus to about 250° out of con-
tact with the air. It is not poisonous, is insoluble in carbon disul-
fide, and does not ignite until heated to 256°. It is non-luminous
in the dark, does not oxidize in the air, but is readily oxidized by nitric
acid to phosphoric acid.

Light-red phosphorus is obtained, according to Schenk,{ by heat-
ing a solution of white phosphorus in phosphorus tribromide for
hours with a return-flow condenser. The phosphorus which then
separates is of a light-red color, is not poisonous, but enters into reac-
tion so readily that it is easily distinguished from red phosphorus.
It dissolves in concentrated potassium or sodium hydroxide with a
stormy evolution of phosphine, the reaction taking place even more
readily than with white phosphorus. When covered with ammonia,
it blackens.

Black phosphorus is obtained when red phosphorus and lead are
heated together in a sealed tube to a red heat and the mass treated
with dilute nitric acid after it is cold; the lead dissolves and leaves
the phosphorus as black phosphorus. By heating to 360° it is changed
to ordinary phosphorus again.

Phosphorus is found in a great many organic substances. In
order to detect its presence, the compound is heated in a sealed tube
with fuming nitric acid, which destroys the organic matter and oxidizes
the phosphorus to phosphoric acid (detected by any of the above
reactions).

Arsenious, arsenic, and chromic acids, which also belong to this
group, have already been described on pp, 225, 229 and 135,

Mitscherlich Test for White, Poisonous Phosphorus {

This sensitive test is based upon the luminescence of white phosphorus when
exposed to moist air in the dark. It is used to detect phosphorus in cases of
poisoning. .

Procedure.—Place the food residues, or finely-cut pieces of the body, in the
liter flask K (Fig. 26) and add enough water to form a thin paste. Then, while
shaking, add tartaric acid to slightly acid reaction, in order to combine with

* Ann. Chem. Phys., 112 (1859), 214.
+ Ber., 36, 979 (1903).
tJ. pr. Chem., 66, 238 (1855).

384 REACTIONS OF THE ACID CONSTITUENTS

any ammonia present. Connect the tube R with the flask and heat the contents
of the latter to boiling. Carry out this operation in a dark room. As the
vapors condense in the tube A, a greenish luminous zone is visible even when only
a few milligrams of phosphorus are present. If larger quantities of phosphorus
are at hand, the distillate in the flask B contains tiny globules of phosphorus
which, by gently heating and rotating the liquid, can be made to collect into a
larger drop; the aqueous solution also contains phosphorous acid which can be
detected by the method of Blondlot-Dusart (see below).

If, therefore, the luminosity is apparent during the distillation in the dark,
then the presence of white phosphorus is probable but not certain, because

Fig. 26.

phosphorus subsulfide, P.8;, sometimes used as a substitute for phosphorus in
the manufacture of matches, will often cause luminescence in the Mitscherlich
apparatus,* especially if a little zine oxide is added to the liquid in K to com-
bine with H.S, which tends to prevent the luminescence.

If the luminescence is not apparent, it is not certain that white phosphorus

*T. Mar and F. Scuarrer, Ber., 1903, 870; L. Vianon, Bull. soc. chem. [3], 33,
805 (1905), and ScnenKk and Scuarrr, Ber., 1906, 1522. The author wishes to
state that all commercial preparations of P,S; do not show luminescence in the
Mitscherlich apparatus. Thus a sample from Kahlbaum did not show the slightest
luminescence by boiling with water or with concentrated salt solution. It was
very pure and contained 55,82 per cent P and 44,14 per cent 8,

PHOSPHORUS 885

is absent; traces of ammonia, hydrogen sulfide, alcohol, ethereal oils and un-
saturated hydrocarbons, interfere with the test. In such cases it is advisable
not to stop distilling too soon, as it often happens that the interfering substances
will distill over and then the luminescence will appear. In case no: luminescence
is noticed, examine the distillate. Treat a part of it with strong chlorine water,
evaporate to small volume on the water-bath and then test for phosphoric
acid. Cf. pp. 378 and 379.

According to J. Peset,* the luminosity is very distinct in the flask if the liquid
is heated to boiling, allowed to cool somewhat, and again boiled. In this way
0.004 mgm. of phosphorus can be detected.

Detection of Phosphorus and Phosphorous Acid according to Blondlot-
Dusart

This beautiful method is based upon Dusart’s observation that hydrogen
containing phosphine, when allowed to flow from a tube provided with a plat-

TANS -——.
Li
enP IOS ‘ )
: |
|
|
|

ouecbaeaeiad

Seeahy

aha

a8
Lx) A
Y)

| TEPPER AHHH 11111

Fia. 27.

inum tip, will burn with a flame having an emerald-green core. The green color
is particularly apparent upon holding a cold porcelain dish in the flame.

Since phosphorus, phosphorous acid and hypophosphorous acid (not phos-
phorie acid) are easily reduced to phosphine by zinc and dilute sulfuric acid,
it is merely necessary to pass the evolved gas through a tube with a platinum
tip, and light it when the air is all expelled; the merest trace of phosphorus is
recognized by the green color.

* Z. anal. Chem., 48, 35 (1909).
+ L. Dusart, Compt. rend., 48, 1126 (1856), and BLuonptot, J. pharm. chim.
(3| 40, 25 (1854),

386 REACTIONS OF THE ACID CONSTITUENTS

Inasmuch as organic substances can prevent the appearance of the green
flame, the phosphorus is first separated from it as follows: Place the solution
containing the phosphorus, or the distillate obtained by the Mitscherlich test,
in a gas-evolution flask, add zine (free from phosphorus) and dilute sulfuric acid
(1 : 7), and pass the evolved gas into a neutral solution of silver nitrate; ~
if phosphorus is present a black precipitate of silver phosphide is obtained,
which, if hydrogen sulfide was present, may contain silver sulfide. Filter off
this precipitate, wash it well with water, and place it in the Blondlot apparatus
(Fig. 27). In the 500 cc. Woulfe bottle, W, generate hydrogen by means of
zinc, free from phosphorus, and dilute sulfuric acid (1: 7). After the air is
entirely expelled from the apparatus, close the pinch-cock a, which causes the
acid to rise into the reservoir 7' (a bottle with the bottom cut off). Now open
the cock a wide enough to permit a steady stream of hydrogen to pass out
from the delivery tube, which is made of potash-glass and is provided with a
platinum tip.* The flame from the lighted gas should not be too large.

If the flame shows no green luminescence in the dark when a porcelain dish
is held in it, then the hydrogen gas is free from phosphorus and can be used —
for the test. Rinse the black silver precipitate through 7' into the bottle W.
If the precipitate contained phosphorus, the core of the flame becomes green,
particularly noticeable upon holding a porcelain dish in it. Any hydrogen sul-
fide evolved collects in the U-tube, U, which contains pumice wet with con-
centrated caustic potash solution.

Since both the Mitscherlich and the Blondlot-Dusart tests give indications
not only of white phosphorus, but also of phosphorus subsulfide, it was desirable
to have a test to serve for the identification of white phosphorus with certainty
even when the sulfide is also present. For this purpose, R. Schenk and E.
Scharfff make use of the property that white phosphorus has of ionizing the
atmosphere, a property which the sulfide does not possess. They use the
Elster-Geitel apparatus for this purpose. For details of the test the original
paper must be consulted.

* A small blowpipe tip can be used here or, still better, a cylinder made by
rolling together some platinum foil.

+R. Scuenk and E. Scuarrr, Ber., 1906, 1522. For the detection of white
phosphorus in the presence of hypophosphorous and arsenious acids see A.
LecuivE, Chem. Zentr., 1912, I, 684,

THIOSULFURIC ACID, H2S203

This very unstable acid, in which one atom of sulfur has a positive
valence of six and the other a negative valence of two, is soon decom-
posed, even in dilute aqueous solution, into sulfurous acid and sulfur:

HeS203 = H2S03+S.

If the aqueous solution of a thiosulfate is treated with dilute hydro-
chloric or sulfuric acid, the solution remains clear for a short time;
but it soon becomes turbid, owing to the deposition of sulfur, which
in this case (unlike most precipitated sulfur) appears yellow.

The salts of thiosulfuric acid, the thiosulfates, are much more
stable than the free acid.

Formation of Thiosulfates
1. By boiling sulfur with an alkali or alkaline-earth hydroxide:
4S+60H- — 28-+8.0;-+3H.,0.

This reaction is analogous to the action of the halogens and of phosphorus
upon hydroxides, forming chloride and hypochlorite, phosphide (phosphine)
and hypophosphite, etc. (cf. pp. 292 and 345).

2. By boiling sulfites with sulfur:

SO;-+S — 8.037.
3. By treating alkali polysulfides with alkali sulfite in the cold:
Na,S;+4Na,SO; =4Na,8,0,+Na:8.
4, By the oxidation of polysulfides:
2NaS2+302 =2Na.8.03.

This last reaction takes place on boiling the solution of polysulfide in the
air, or very slowly on standing. Yellow ammonium polysulfide is changed, on
standing in the air, into ammonium thiosulfate with deposition of sulfur.

The sulfites can be kept well in aqueous solutions, provided they are not
subjected to the action of carbon dioxide. They are gradually decomposed by
the latter, with separation of sulfur.

The most important commercial thiosulfate is the sodium salt NasS.0;-5H,0,
the well-known ‘ hypo ” of photographers.

Solubility.—The thiosulfates of the alkalies are readily soluble
in water, the remaining ones are difficultly soluble; many of them
dissolve in an excess of sodium thiosulfate, forming complex ions.

387

388 REACTIONS OF THE ACID CONSTITUENTS

REACTIONS IN THE WET WAY

Use a solution of sodium thiosulfate for these reactions.

1. Sulfuric Acid.—Both dilute and concentrated sulfuric acid
decompose thiosulfates, with deposition of yellow sulfur.

2. Silver Nitrate produces a white precipitate, which rapidly
becomes yellow, brown, and finally black, owing to the formation of
silver sulfide:

S2037+2Agt — AgoS203;
AgoS203 +H:2O = HeS04 +AgoS.

Silver thiosulfate is soluble in an excess of the reagent. Diffi-
cultly soluble Na[AgS2Os3] is at first formed, ’

Ag2S203+8203° — 2[AgS20s!,

which combines with more thiosulfate, forming a soluble complex
salt:

2[AgS203]"+8203— — [Age(S203)3])" ~.
But by boiling the dilute solution, silver sulfide is precipitated:

Ago(S203)3]~ ~ > S2037 +804" +802 T +S+AgeS.

Many other metals behave like silver, especially those of the
hydrogen sulfide group. Thus copper, mercurous, and tin salts are
precipitated as sulfides by boiling the acid solutions with sodium
thiosulfate.*

3. Barium Chloride in excess produces a white, crystalline pre-
cipitate of barium thiosulfate,t difficultly soluble in cold water (500
cc. of water at 18° dissolve about 1 gm. of BaS2O3), but fairly soluble
in hot water.

4. Strontium Chloride produces a white, crystalline precipitate,
but only in very concentrated solutions (3.7 cc. of water at 18° dis-
solve 1 gm. of SrS2Os).

5. Lead Acetate precipitates white lead thiosulfate, solulite in an
excess of the alkali thiosulfate. On boiling the solution a volumi-
nous precipitate, consisting of lead sulfate and lead sulfide, is formed.

6. Iodine Solution is decolorized by a thiosulfate solution:

28203 +12 — 21°-+840¢6-.

* Z. anorg. Chem., 28, 223 (1902). S
{ Rubbing the sides of the test-tube hastens the formation of this precipitate.

THIOSULFURIC ACID 389

The iodine is reduced to iodine anions and the thiosulfate ion is oxidized
to tetrathionate ion.
Chlorine and bromine in excess (cf. D. 305) act quite differently
upon thiosulfates. If chlorine (or bromine) is conducted into a solu-
tion of sodium thiosulfate, a considerable precipitation of sulfur takes
place, which, upon further action of the halogen, disappears:

S203° +He0+Cle — 2Cl>+ 2HtS04-+S;
S+4H.0+3Cle=8H*t+6Cl- +804".

Other weak oxidizing agents act in the same way as iodine. Thus,

7. Ferric Chloride produces, in’ solutions of sodium thiosulfate,
at first a dark-violet coloration (perhaps ferric thiosulfate), which
disappears after some time, leaving a colorless solution containing
ferrous chloride and sodium tetrathionate:

28203° + 2Fet + 33 = 2Fet t48406.
Similarly,
8. Cupric Salts are reduced to colorless cuprous compounds, with
the formation of sodium tetrathionate:

25203 + 2Cut + => Cut T4. S406 .

The unstable cuprous sulfate immediately acts upon more thio-
sulfate, forming sodium cuprous thiosulfate:

Ug? *+4+280037 — [Cue(S203)2]>.

If the colorless solution of the cuprous salt is treated with caustic potash
solution, yellow cuprous hydroxide is in some cases immediately formed, in
other cases only on standing or on warming. The precipitate becomes darker
colored on being boiled.

If the solution is acidified and boiled, black cuprous sulfide is precipitated.

The colorless solution of the cuprous salt also gives a white (usually a light
pink) precipitate with potassium ferrocyanide or cuprous ferrocyanide.

9. Nascent Hydrogen (zinc and hydrochloric acid) causes the
evolution of hydrogen sulfide.

10. Zinc Salts produce no precipitate (difference from sulfides).

11. Zine Sulfate and Sodium Nitroprusside produce no red colora-
ation (difference from sulfites).

12. Potassium Cyanide.—Boiling a solution of a thiosulfate with
potassium cyanide and caustic soda transforms the thiosulfate into aut
fite and the cyanide into thiocyanate:

$203--++-CN~ — SO3"=+CNS~

390 REACTIONS OF THE ACID CONSTITUENTS

On acidifying the solution with hydrochloric acid and adding
ferric chloride, the blood-red color of ferric thiocyanate is obtained
(difference from sulfites).

Detection of Sulfurous and Thiosulfuric Acids in the Presence of
Hydrogen Sulfide

A Method of E. Votocek.*—Principle.—Alkali sulfites, sulfides and polysul-

fides in slightly alkaline solution will decolorize fuchsin, malachite green or a
mixture of these two dyestuffs. Ifa solution of acetaldehyde or of formalin is
added to the decolorized solution, the color returns. Sulfhydrates, thiosulfates
and thionates do not decolorize a solution of the above dyestuffs.

Reagent.—Dissolve 0.025 gm. of fuchsin and 0.025 gm. of malachite green
separately in 100 cc. portions of water. Mix three volumes of the fuchsin
solution with one volume of the malachite green solution.

Procedure.—It is assumed that the solution is slightly alkaline. Test the
solution first for sulfide (monosulfide, sulfhydrate and polysulfide) by treating
a little of it with 2 or 3 drops of sodium nitroprusside solution. A reddish-
violet color shows the presence of the sulfide anion. If sulfide is present, treat
the remainder of the solution with cadmium carbonate, shake vigorously, and
allow the cadmium sulfide to settle somewhat. Filter and test a new portion
of the filtrate with sodium nitroprusside to see if all of the sulfide has been
removed. When all the sulfide has been removed or proved absent, treat the
remainder of the filtrate with a drop of phenolphthalein solution and intro-
duce carbon dioxide gas until the solution is decolorized by it. Take 2 or 3 ce.
of this colorless solution and test it with 2 or 3 drops of the fuchsin-malachite-
green reagent. If the color solution is decolorized, a sulfite is present. To the
remainder of the solution add a little dilute hydrochloric acid, boil a few minutes,
and notice whether there is any deposition of sulfur. If the solution remains
clear, no thiosulfate is present.

This is the best method for detecting a sulfide, a sulfite and a thiosulfate |

in the presence of one another.

(b) Method of Autenrieth and Windaus.}—The three acids are assumed
to be present together in solution in the form of their alkali salts. Treat the
fairly concentrated solution with cadmium carbonate, shake and filter off the
excess cadmium carbonate and any cadmium sulfide which will be formed if
a sulfide is present. Treat the filtrate with strontium nitrate solution and
allow it to stand overnight. Filter off any strontium sulfite that may be
formed and wash it with a little cold water. If the strontium sulfite is treated
on the filter with dilute hydrochloric acid, sulfurous acid goes into solution,
which can be detected by its property of decolorizing an iodine solution. In
the filtrate from the strontium sulfite, the thiosulfate remains; it can be

detected by acidifying with hydrochloric acid and warming, when sulfur will

be deposited,

* Ber., 40, 414 (1907). |
+ Z. anal. chem., 1898, 295. For another method of detecting sulfite in the
presence of thiosulfate, cf. F. E. Weston, Chem. Zentr., 1910, I, 379.

.
OU

THIOSULFURIC ACID 391

Solubility of Sulfites and Thiosulfates of the Alkaline Earths in Water.

Sulfite. Thiosulfate.
SERSOMIR I OR aah eis ake: 1 : 800 Fee
Strontium..... relate PS Sra ean 1 : 30,000 LaF
ESCM A rear anak eee ae 1 : 46,000 1 : 480

REACTIONS IN THE DRY WAY

The thiosulfites of the alkalies, on being heated out of contact
with the air, are changed into sulfate and polysulfide, and the latter
into sulfide and sulfur:

4Na2Se03 — 3Na2SO4+NasSs,
Nass ae Naes +48.

If this reaction is performed in a closed tube, a sublimate of sul-
fur is obtained (difference from sulfites); and the residue yields hydro-
gen sulfide if treated with acid.

GROUP V

Silver Nitrate produces no precipitate in acid or neutral solutions.
Barium Chloride, also, causes no precipitation.

NITRIC ACID, HNO;

Occurrence.—Nitric acid is found in the form of nitrates in small
amounts almost everywhere in nature; thus the ammonium salt
is found in the atmosphere and in soils; the calcium salt is found
in old masonry; while the sodium salt is found in rainless localities,
particularly in Chili (Chili saltpetre).

Nitric acid is the final product of the oxidation of ammonia; it is
found wherever nitrogenous organic substances have been subjected
to decay, forming ammonia. With the help of microérganisms
(Monas nitrificans, according to Winogradsky) the ammonia is changed
first to nitrous acid,

2NH3+302 = 2H20+2HNOdzg,
and by further oxidation to nitric acid:
2HNO2+02 =2HNOs.

Properties—Pure nitric acid is a colorless liquid, with a specific
gravity of 1.54 at 20°. At 86° it begins to boil, with decomposition,
giving off its anhydride, which suffers further decomposition into
nitrogen peroxide, NOz (brown fumes), and oxygen. By the constant
loss of N2Os, the nitric acid becomes more and more dilute and the ©
boiling-point constantly rises, until at 120.5° C. it remains constant;
when nitric acid of specific gravity 1.414 distills over, forming a 68
per cent acid. Ifa more dilute acid is subjected to distillation, water
is at first given off, the boiling-point constantly rising until 120.5° C.
is reached, when a 68 per cent acid again distills unchanged.

Red, fuming nitric acid is obtained by conducting NOe into the
colorless, concentrated acid. In its most concentrated condition it
possesses a specific gravity of 1.55.

If the fuming acid is treated with water, it is colored green, and

vapors of nitric oxide are given off, which are colored brown on com-
392

NITRIC ACID 393

ing in contact with the air. The dissolved NOz (or, better, N2O,),
being a mixed anhydride, is changed into nitric and nitrous acids,

N204+H20 — HNOs+HNOzg,

and the nitrous acid, owing to the heat of reaction, is partly changed
into nitric acid, with evolution of nitric oxide,

: 3HNO2=H20+HNO3+2NO 7 ,
and
NO+0O=N0Oz 7 (brown vapors).

Nitric acid is a strong oxidizing agent (cf. p. 30). It is mono-
basic, and, next to the halogen acids, is the strongest acid (cf, p. 10).
It forms stable salts, which are all soluble in water; but a few of them
are changed by water into basic salts (cf. bismuth and mercuric salts),
insoluble in water, but soluble in dilute nitric acid.

REACTIONS IN THE WET WAY

As nitric acid does not form insoluble salts, it cannot be detected by means
of precipitation; its characteristic reactions depend upon its oxidizing action.
Great care must be exercised before deciding whether this acid is present, for
other oxidizing substances give similar (in some cases the same) reactions.

1. Dilute Sulfuric Acid gives no reaction (difference from nitrous
acid).

2. Concentrated Sulfuric Acid when heated with any nitrate causes
evolution of yellow to brown vapors of NOs, with a characteristic
penetrating odor.

3. Silver Nitrate and Barium Chloride cause no precipitation.

4. Ferrous Salts are oxidized by nitric acid, which is itself reduced
to nitric oxide, NO. |

If the reaction takes place in the cold, the latter combines with the excess
of ferrous salt, forming a dark-brown, very unstable compound, FeX,-NO.
This compound is decomposed, on warming, into ferrous salt and nitric oxide
(which escapes) the brown color disappearing. If the amount of nitric acid
present is more than sufficient to oxidize completely the ferrous salt to ferric
salt, a more reddish coloration is obtained owing to the formation of a salt such
as Fes(SO4)3- 4NO.

The oxidation of the ferrous salt takes place according to the following

reaction,
3Fett+NO; +4Ht — 3Fe++++2H,0+N0O T

and is best carried out in the following manner:
Place some of the substance to be tested for nitrate in a test-tube with 5 ce.
~ of water, add an equal volume of concentrated sulfuric acid and cool nearly toe

394 REACTIONS OF THE ACID CONSTITUENTS

room temperature. by shaking the solution under running water. Carefully
add about 5 ec. of saturated ferrous sulfate solution down the sides of the test-
tube so that it forms a layer on top of the concentrated sulfuric acid solution.
When a nitrate is present a brown ring is formed at the zone of contact between
the ferrous sulfate solution and the heavier sulfuric acid solution, or if con-
siderable nitric acid is present the whole of the ferrous sulfate solution may be
colored. If only a little nitric acid is present, the zone may be colored pink
owing to the formation of Fe.(SO,.);-4NO. The test is not reliable in the pres-
ence of an iodide or a chromate.

Nitrous acid gives the same reaction, with the difference that
it takes place even without the addition of concentrated sulfuric acid.

5. Indigo Solution is decolorized by warming with nitrie acid
(as well as by other oxidizing agents). eee

6. Potassium Iodide is not decomposed by pure, dilute nitric
acid (difference from nitrous acid).

If to the solution of a nitrate we add potassium iodide, a few drops of an
acid (best acetic acid), and a little zinc, the nitric acid is reduced to nitrous
acid, which then reacts with hydriodic acid so that the solution becomes yellow
on account of the separation of iodine. By shaking the solution with carbon
disulfide, the latter will be colored reddish violet, or the iodine may be
detected by adding a little starch paste.

The reactions which take place may be represented by the following equa-
tions:

Zn+NO;-+2Ht — Znt++NO,.-+H.0;

2NO,_+2I1-+4H* — 2NO ft +2H,0+1.

If it is desired to detect the free iodine by forming a solution in carbon
disulfide, the latter wnder no circumstances should be added before the zine
-has been allowed to act upon the acid solution of the nitrate and potassium
iodide. In such a case, there will often be no separation of iodine, because
the nascent hydrogen is used up, reducing the carbon disulfide to thioformalde-
hyde and hydrogen sulfide, and the latter reacts with any iodine which may be
formed, changing it back to hydriodic acid:

CS.+2H:, = CH.S +H,§;
H.S+2I =2HI+S8.

7. Diphenylamine Reaction (the Lunge test *).—Reagent.—Dissolve
0.5 gm. of diphenylamine in 100 cc. of pure, concentrated sulfuric acid
diluted with 20 cc. of water.

Procedure.—Place a few cubic centimeters of the diphenylamine solution
in a test-tube and carefully cover it with the solution to be tested for nitric
acid. If the latter is present, there is formed at the zone of contact between the
two liquids a ring of a beautiful blue color.

* LunGE, Z. angew. Chemie, 1894, 345.

NITRIC ACID 395

This very sensitive reaction is, unfortunately, also caused by nitrous,
chloric, and selenic acids, ferric chloride, and many other oxidizing agents.
Even fuming sulfuric acid will give it sometimes.

In the absence of ferric and selenic salts, it is useful for detecting the presence
of small amounts of nitrogen acids in sulfuric acid. In this case first pour the
concentrated sulfuric acid to be tested into the test-tube and cover it with the
specifically lighter diphenylamine solution. If 1 cc. of an acid containing
only 3; milligram of nitrogen in a liter is used, the reaction will cause a notice-
able coloration. If very strong fuming sulfuric acid is to be tested, dilute it first
with concentrated sulfuric acid until it does not contain more than 20 per cent
excess SOs. 3

8. Brucine Reaction.—Reagent.—Dissolve 0.2 gm. of brucine in
100 cc. of pure concentrated sulfuric acid.*

Procedure.—Mix the solution to be tested for nitric acid with three times its
- volume of pure, concentrated sulfuric acid, and add 1 ce. of brucine solution.
If nitric acid is present, a red coloration quickly appears, which quickly changes
to orange, then slowly to lemon or gold yellow, and finally becomes greenish
yellow. Nitrous acid does not give this reaction provided it is present as
“ nitrose,” 7.e., dissolved in concentrated sulfuric acid. Aqueous solutions of
nitrites always yield a small amount of nitric acid when acidified with sulfuric
acid, and consequently give the brucine reaction.

9. Zinc in Alkaline Solution reduces nitric acid to ammonia.
If a nitrate solution is boiled with zinc dust and an alkali, a con-
siderable evolution of ammonia takes place:

NO3~+4Zn+70H- — 4Zn02=+2H20+NHs } .

Devarda’s alloy reacts much more quickly with a drop of caustic
soda. This reaction is particularly suited for the detection of nitric
acid in the presence of chloric acid (cf. p. 399).

Detection of Nitric Acid in the Presence of Nitrous Acid

With the exception of the Lunge-Lwoff method, there is no absolutely
reliable qualitative test for the detection of traces of nitric acid in the presence
of large amounts of nitrous acid in aqueous solution. A number of methods
have been proposed which depend upon the destruction of the nitrous acid by
diazotizing, but they all yield only approximate results; because, in order to
destroy the nitrous acid, it is necessary first to set the acid itself free by the
addition of another acid, which always causes a part of the nitrous acid to be
changed to nitric acid; so that the latter will be detected even when no nitric
acid was originally present.

Large amounts of nitric acid in the presence of nitrous acid may be detected
by the method proposed by Piccini,} in which a concentrated solution containing

* Lunan, Z. angew. Chemie, 1894, 348.
{ Z. anal. Chem., 19, 354.

396 {REACTIONS OF THE ACID CONSTITUENTS

salts of both acids is treated with a concentrated solution of urea, and then
covered (by means of a pipette) with dilute sulfuric acid. A lively evolution of
nitrogen and carbon dioxide ensues, which ceases in a few minutes:

CO(NE,): +2HNO, =CO, T +38H,0+2N, 1.
Urea
When the evolution of gas has ceased, test the solution for nitric acid by
means of the diphenylamine reaction.
This reaction, however, does not take place quickly enough to prevent
traces of nitric acid being formed according to the following equation: .

3HNO. =H.0+HNO;+2NO 1.

The odor of nitrous fumes is always perceptible in the escaping nitrogen,*
which can also usually be detected by means of iodo-starch paper. The nitric
acid which remains in the solution can be detected by means of the diphenyl-
amine reaction.

The nitrous acid may also be destroyed by boiling an alkaline nitrite solu-—
tion with neutral ammonium chloride; but traces of nitric acid are always in
evidence at the same time.f

If the diphenylamine reaction gives a very intense coloration after the
destruction of the nitrous acid by means of urea, the presence of nitric acid in
the original compound is assured; but if the reaction shows that only a trace of
nitric acid is present, it is probably due simply to small amounts of nitric acid
formed by the destruction of the nitrous acid.

REACTIONS IN THE DRY WAY

By the ignition of nitrates of the alkalies, they are changed into |
nitrites with loss of oxygen, and the latter are decomposed on stronger
ignition into oxide:

2KNO3=2KNO2+Ob2 fT,
4AKNO2=2K20+4N0+02 7.

All nitrates deflagrate on being heated on charcoal; i.e., the char-
coal burns at the expense of the oxygen of the nitric acid, witli vivid
scintillation. ;

* Even at 0° and in an atmosphere of carbon dioxide.
} By evaporating with ammonium carbonate solution the decomposition scarcely
takes place at all.

CHLORIC ACID, HC103

Free chloric acid is extremely unstable, and is decomposed at
40°, into perchloric acid with loss of chlorine and oxygen (cf. p. 170):

3HC103 > H20+2C102 T +HC10,,
and
2Cl02=Cle T +202 Tf.

The salts of the monobasic chloric acid, the chlorates, are quite stable
and are all soluble in water. They are formed by conducting chlorine
into hot alkali hydroxide solutions, which are not too concentrated.*

60H +38Cle — 5CI-+ClOs3"+3H20.

REACTION IN THE WET WAY

1. Dilute Sulfuric Acid sets free chloric acid from chlorates, which,
as above stated, is gradually decomposed, with loss of chlorine and
oxygen, into perchloric acid. The solution, therefore, acts as an
oxidizing agent, particularly on warming; it will turn iodo starch
blue:

ClO3-+61-+6H* > Cl-+3H20+4+3lz.

The speed of the reaction between the free chloric acid depends upon
the concentration of the hydrogen ion. The neutral salts do not
- act as oxidizing agents (difference from hypochlorites).

Test for Hypochlorite in Chlorates.

To test for the presence of hypochlorite in an alkali chlorate, dissolve about
2 gm. of the salt in 200 cc. of water, add 3 cc. of 10 per cent KI solution and 3 ce.
of starch solution, but no sulfuric acid; the solution will at once turn blue if a
trace of hypochlorite is present. As little as 0.1 mgm. of hypochlorite will
give the test.

2. Concentrated Sulfuric Acid decomposes all chlorates, setting
free greenish-yellow chlorine dioxide gas, which explodes violently
on warming:

8KC103+3H2SO4 = 3KHS04+HC101+2C102 T +H20.

*F, WINTELER, Z. anorg. Chem., 38, 188 (1902).
397

398 REACTIONS OF THE ACID CONSTITUENTS

3. Silver Nitrate and Barium Chloride do not cause precipitation.
4. Reducing Agents reduce chlorates to chlorides in acid, alkaline,
and neutral solutions.

The reduction in acid solution is effected by means of zine and dilute sul-
furic acid, or by means of sulfurous acid:

3Zn+ClO;-+6Ht —> 8Zn+++Cl-+3H,0;
3S0;-+Cl0;- > Cl-+3S0--.

The reduction in alkaline or neutral* solution is brought about by boiling
the solution with zinc dust, or better, by means of Devarda’s alloy (cf. p. 34):

ClO;-+3Zn+60H~ — 3Zn0.-+Cl-+3H.0.

The residue of zinc dust (or copper, if the alloy is used) is filtered off, the
solution acidified with nitric acid,f and silver nitrate added, when the charac-
teristic, curdy precipitate of silver chloride is formed.

5. Concentrated Hydrochloric Acid decomposes all chlorates, with
evolution of chlorine:

ClO3~--+5CI-+6H* > 3H20+3Ck f .

This equation is not correct, for some ClO, is always mixed with the Ch.
The following equation expresses this:

3Cl0;-+7Cl-+10Ht — 5H.0+4Cl f +2Cl0, ft.

The proportion of ClO, and Cl, formed is influenced by the concentration of
the reacting substances and the temperature.

6. Ferrous Salts.—By boiling chlorates with ferrous salts in the
presence of dilute acid, the chlorate is quickly reduced to chloride
(difference from perchloric acid): .

ClO3-+6Fe* *+6H* — 3H20 +Cl"+6Fe* aS

7. Diphenylamine reacts the same as with nitric acid.

Detection of Hydrochloric, Nitric, and Chloric Acids in the Presence
of One Another

I. First, test for the presence of chlorine anions by treating a part of the
solution with silver nitrate; a white precipitate of silver chloride shows the pres-
ence of hydrochloric acid. Treat the remainder of the solution with silver
sulfate solution until no further precipitation of silver chloride takes place, and
filter off the precipitate.

* The reaction takes place very slowly in neutral solutions.
7 On acidifying with nitric acid a heavy precipitate of Zn(OH)2 is obtained
which dissolves in more nitric acid.

CHLORIC ACID 399

Boil the filtrate with a little caustic potash (in order to expel any ammonia
from ammonium salts which may be present), add zine dust (or Devarda’s
alloy), and again boil; if nitric acid is present, ammonia will be given off. Filter
off the residue, acidify the filtrate with nitric acid and treat with silver nitrate.
If a precipitate of silver chloride is now obtained, chloric acid was originally
present.

II. Or, test a small part of the solution for hydrochloric acid by adding silver
nitrate in excess, filter off the precipitate, treat the filtrate with sulfurous acid,
and again test with silver nitrate: a precipitate of silver chloride shows the
presence of chloric acid. Enough nitric acid must be added to dissolve any
silver sulfite that may form.

Test a second portion of the solution, as above, for nitric acid.

REACTIONS IN THE DRY WAY

On ignition, all chlorates are decomposed, forming a chloride
with loss of oxygen. By heating on charcoal, deflagration takes
place.

PERCHLORIC ACID, HClO,

Free perchloric acid is obtained by the distillation of potassium
perchlorate with concentrated sulfuric acid. In this way the solid,
crystalline- hydrate HClOs+H20O is obtained, and, on heating to
110°, the anhydrous liquid acid distills off first and fumes strongly
in the air, while the oily hydrate HClOs4+2H20 remains behind until
the temperature reaches 203° C., when it also distills.

The concentrated acid is very dangerous, and often explodes spon-
taneously. In aqueous solution, however, it can be kept without
danger. 7

The salts of this monobasic acid, the perchlorates, are remarkably
stable; they contain chlorine with seven positive charges and are
isomorphous with the permanganates. The potassium salt is obtained
from potassium chlorate. On melting the latter compound, at first
a lively stream of oxygen is given off which, however, soon lessens
The melt quickly becomes viscous, and consists of potassium chloride
and potassium perchlorate,

2KCIO3 = KCI4+ KClO4+0Oc2 7 ,

and the latter may be separated from the much more soluble potassium
chloride by recrystallization. ,
Solubility —All perchlorates are soluble in water.

REACTIONS IN THE WET WAY

Perchloric acid is not attacked by concentrated sulfuric acid,
nor reduced to chloride by zinc dust, Devarda’s alloy, sulfurous acid,
or acid solutions of ferrous salts. .

1. Potassium Salts precipitate the relatively insoluble, white,
erystalline KClOs (cf. p. 81). :

2. Silver Nitrate and Barium Chloride produce no precipitation.

REACTIONS IN THE DRY WAY

The perchlorates deflagrate on being heated on-charcoal; by fusing
they lose oxygen, leaving chloride behind, which when dissolved in
water gives all the reactions for hydrochloric acid.

400

PERSULFURIC ACID, H2S20g

Pure persulfurie acid itself has never been isolated, its solution
in sulfuric acid alone being known. It was first prepared by M.
Marshall,* who electrolyzed fairly dilute sulfuric acid, keeping it
very cold. During the electrolysis hydrogen ions are discharged
at the cathode and unite to form hydrogen molecules, while HSO4
anions are discharged at the anode and unite to form persulfuric acid:

2H2SO4 > He T +H2820s.

The preparation of ammonium persulfate, from which all other
persulfates are made, is entirely analogous.

The most important salts of persulfuric acid are those of ammo-
nium, potassium, and barium. (NH4)2SeOg is readily soluble in
water, and forms monoclinic crystals; K2SeOg is difficultly soluble
in cold water, but much more soluble in hot water, from which solu-
tion it is obtained by rapid cooling in the form of long crystals;
BaS20g+4H20 is made by rubbing ammonium persulfate with barium
hydroxide, and is fairly soluble in water.

REACTIONS

A solution of ammonium persulfate may be used.

1. Water.—All persulfates are decomposed in aqueous solution
(slowly in the cold, but more quickly on warming), forming sulfate,
free sulfuric acid, and oxygen:

2520s +2H20 — 4HSO04°- +02 fT ,
2BaS20s+ 2H20 — 2BaS0O.+2H2S04+02 t e

A large proportion of the oxygen escapes as ozone, which can be
detected by its odor, or by its property of turning iodo-starch paper
blue. A dilute solution of ammonium persulfate decomposes slowly
at 20° C., without evolution of oxygen, part of the nitrogen being.
oxidized to nitric acid:

8(N H4) 2820s-+6H20 = 14(NH4)HS804+2H2S04+2HNOs.

* J. Chem. Soc., 59, 771.

401

402 REACTIONS OF THE ACID CONSTITUENTS

2. Dilute Sulfuric Acid acts the same as water.

3. Concentrated Sulfuric Acid.—If a solid: persulfate is dissolved
in concentrated sulfuric acid at 0° C., a liquid is obtained which pos-
sesses very strong oxidizing properties. The mixture is known as
Caro’s acid.* For further particles concerning this acid see p. 403. —

4. Silver Nitrate precipitates black silver peroxide:

2Agt+ S203 +2H20 — 2HSO4 + 2H*+AgoOo.

If, however, the concentrated solution of ammonium persulfate is —
treated with ammonia and a very little silver nitrate, a lively evolution
of nitrogen takes place, and the solution becomes heated to boiling.
Silver peroxide is formed first, and oxidizes the ammonia to water,
setting free nitrogen (catalysis).f

5. Manganese, Cobalt, Nickel, and Lead Salts are oxidized in
the presence of alkali to black peroxides:

Mntt+ SeOs +40H — 2804>+H2eMn03+H20.

In this last reaction persulfuric acid reacts exactly similarly to hydrogen
peroxide. It may be distinguished, however, from the latter by the fact that
it does not decolorize a solution of potassium permanganate, does not produce
a yellow coloration with titanium sulfate, and does not react with chromie acid
to form chromium peroxide (cf. p. 139). Ferrous salts are readily oxidized
to ferric salts, and cerous salts are changed to yellow ceric salts by persulfates,
but the latter are not decolorized by an excess of the persulfate, while they are —
by hydrogen peroxide.

Manganese and lead salts are precipitated quantitatively from neutral and
slightly acid solutions by alkali persulfates, cobalt incompletely from neutral
solutions and not at all from acid ones, and nickel only in the presence of alkali.
Hydrogen peroxide produces precipitates of peroxides in all these solutions only
in the presenceof alkali. In the presence of silver ions, which have a catalytic
effect, manganous ions are oxidized to permanganate in ‘hot nitric acid solu-
tions by means of alkali persulfates.

6. Barium Chloride does not give a precipitate immediately in
a freshly prepared cold solution of a persulfate; but, on standing
some time, or on boiling, insoluble barium sulfate is precipitated.

* Z. angew. Chem., 1898, 845; Ber., 34, 853 (1901); Ber., 41, 1839 (1909).
+ Z. phys. Chem., 37, 255 (1901).

PERSULFURIC ACID 403

Monopersulfuric Acid (Caro’s Acid) H2:SO;

This acid is formed by the hydrolysis of persulfuric acid,
H.8,0;+H20 — H.S0,+H.S80,,
and by the action of perhydrol (30 per cent hydrogen peroxide) on sulfuric acid:
H.S0,+ H,0: — H.O+H,SO;.

Unlike hydrogen peroxide, it does not reduce permanganates and unlike
persulfuric acid it causes the immediate liberation of iodine from potassium
iodide solutions. |

It is usually assumed that hydrogen peroxide, persulfuric acid and mono-
persulfuric acid all contain an atom of oxygen directly connected to another
atom of oxygen; this atom of oxygen therefore has a positive and a negative
charge residing upon it (cf. p. 33).

GROUP VI

Silver Nitrate produces no precipitate.
Barium Chloride produces a white precipitate, insoluble in acids.

SULFURIC ACID, H2SO,4

Pure sulfuric acid at ordinary temperatures is a colorless, oily
liquid of specific gravity 1.838; at low temperatures it is a solid. If
the acid is subjected to distillation, it is always partially decomposed;
heavy, white vapors of SO3 are given off first, and at 338° C. a 98
per cent acid distills over. Ordinary commercial sulfuric acid has a
specific gravity of 1.83-1.84, and contains 93-96 per cent H2SOx.
It often contains lead sulfate, selenic acid, platinum, palladium, arsen-
ious acid, the nitrogen acids, and small amounts of organic matte
(whereby it is often colored brown) as impurities.

Concentrated sulfuric acid is very hygroscopic, and is used, there-
fore, for drying gases, ete.

The anhydride of sulfuric acid, SO3, dissolves in concentrated
sulfuric acid, forming pyrosulfuric acid, H2S207, which is solid at
ordinary temperatures, melts at 35°, and loses SO3 at higher tempera-
tures. It fumes strongly, and is called, therefore, fuming sulfuric
acid. Sulfuric acid is dibasic and forms both neutral and acid salts.

Solubility — Most: sulfates are soluble in water; calcium sulfate
is difficultly soluble, strontium and lead sulfates are very difficultly
soluble, while barium sulfate is practically insoluble in water. There
are also a number of basic sulfates, Hg, Bi, Cr, which are insoluble
in water, but are, as a rule, easily dissolved by dilute acid.

REACTIONS IN THE WET WAY

1. Sulfuric Acid naturally gives no reaction.

2. Silver Nitrate causes no precipitation in dilute solutions, but
in concentrated solutions a white crystalline precipitate is formed
(100 ec. of water dissolve at 18° C. only 0.58 gm. of silver sulfate).

3. Barium Chloride precipitates, from even the most dilute solu-
tions, white barium sulfate, insoluble in acids. ;

4, Lead Acetate precipitates white Jead sulfate, soluble in con-

404

SULFURIC ACID 405

centrated sulfuric avid, ammonium acetate, and ammonium tartrate
solutions (cf. p. 210).
To detect the presence of SQ, in insoluble sulfates, tréat with sodium
carbonate, whereby insoluble carbonate and soluble sodium sulfate are formed.
Lead sulfate and calcium sulfate are easily decomposed by boiling with
sodium carbonate solution, but barium and strontium sulfates are only incom-
pletely decomposed by this treatment; they are much more readily attacked by
fusing with four times as much sodium carbonate (cf. p. 109).

5. By Nascent Hydrogen (zinc and acid) the sulfates are not
reduced.

REACTIONS IN THE DRY WAY

The neutral salts of the alkalies melt with difficulty without
being decomposed, while the acid salts of the alkalies readily give
off water and SOz3 (cf. p. 130). |

The sulfates of the alkaline earths and of lead do not undergo
decomposition on ordinary ignition; the remaining sulfates are more
or less decomposed.

All sulfates are reduced to sodium sulfide when heated with sodium
carbonate on charcoal; if the product is placed upon a bright silver
coin and moistened, a black stain of silver sulfide results, e.g.::

(a) CaSO4+NazgCOz = CaCO3+ Na280z;
(b) Na2SO4+2C =2COz T +Naz8;
(c) NagS+Age+H20+0 =2Na0H+Ag.S.

This reaction is called the Hepar reaction.

- HYDROFLUORIC ACID, HF

Occurrence.—Hydrofluorie acid occurs in nature only in the form
of fluorides, of which the most important is fluorite, CaF 2, crystalliz-
ing in the isometric system. It is also found as cryolite, Nas[AIFs],
_ In Greenland, and in many silicates, such as tourmaline, topaz, lepido-
lite, apophyllite, apatite, etc.

Preparation.—Hydrofluoric acid is obtained by decomposing
a fluoride with concentrated sulfuric acid in platinum or lead retorts:

CaF2+He2SO4=CaSO4+2HF 7.

Properties—Hydrofluoric acid at temperatures above 20° C. is
a colorless gas which changes at 19.4° C. to a mobile, fuming liquid.
The vapors possess a penetrating odor and are poisonous. When in
contact with the skin, the acid produces painful burns. On heating
the concentrated aqueous solution, the gas HF at first distills and then
the 36 per cent acid. |

Hydrofluoric acid is distinguished from all other acids by its ability to
dissolve silicic acid, a property which is utilized technically for etching glass,
and in the analytical laboratory for detecting fluorine and silicic acid, as well as
for decomposing silicates. On account of this action upon glass the acid must
be kept in platinum, wax, or hard rubber, and prepared in platinum or wick
vessels.

The action of hydrofluoric acid upon silicic acid takes place according &
the equation

SiO. 4HF =2H,0+SiF, 1,

and the velocity of the reaction depends upon the fineness and nature of the
material. Thus Mackintosh * found after one hour’s action of an excess _
of 9 per cent hydrofluoric acid upon quartz and opal powder that the quartz
had lost only 1.56 per cent of its original weight while the opal had lost 77.28
per cent.

If precipitated and ignited silica is treated with strong hydrofluoric acid,
it dissolves almost immediately with hissing and strong evolution of heat,
while quartz powder under the same treatment is dissolved only slowly.

Most silicates, with regard to the ease with which they are attacked by
hydrofluoric acid, stand intermediate between the precipitated silica and quartz,
although some silicates are attacked more difficultly than is quartz, and a few
are only slightly acted upon.

* Chem. News. 64, 102,

406

HYDROFLUORIC ACID 407

Hydrofluoric acid is a weak, monobasic acid having, in common with other
weak acids like carbonic acid, acetic acid, etc., the property of turning blue
litmus red and Brazil-wood paper yellow. The aqueous solution of an alkali
fluoride has a strong alkaline reaction.

The property of forming very. stable complex metal-hydrofluoric acids
is characteristic of hydrofluoric acid as of hydrocyanic acid,

K[Ag(CN),.]:—H[AgF.], H[KF:], H[NaF.] and H[NH.F,],
and corresponding to ferricyanic acid, H;[Fe(CN).]:
Na;[FeF’s], Na,[AIF 6], etc:

Unlike the complex cyanogen compounds, of which the free acids either
do not exist at all or represent very unstable compounds, the corresponding
fluorine acids are fairly stable. Thus hydroargentifluoric acid, H{AgF,]
decomposes only on gently heating it, into silver fluoride and hydrofluoric
acid, and the corresponding alkali compounds are decomposed only upon
ignition; for this reason the latter are suitable for attacking difficultly decom-
posable silicates, zircon and titantium minerals, etc., which are only partially
attacked by free hydrofluoric acid.

Solubility—The fluorides of the alkalies, of silver, aluminium,
tin, and mercury are soluble in water, while those of the alkaline
eniths: of lead, copper, and zinc, are insoluble, or at least very dif-
ficultly soluble.

REACTIONS IN THE WET WAY

For reactions 1, 2, and 3, use powdered calcium fluoride, but for reactions
4, 5, 6, and 7, use a solution of sodium fluoride.

1. Dilute Sulfuric Acid causes only a slight reaction.
2. Concentrated Sulfuric Acid reacts readily on warming, setting
free hydrofluoric acid:

CaF2+H2SO4=CaSO1+2HF }.

Acid containing about 90 per cent H.SO, is most suitable for this reaction.
Acid containing an excess of SO; is likely to cause the formation of difficultly
volatile fluorsulfonic acid, HSO,F.

If the reaction is performed i in a test-tube, the hydrofluoric acid will attack
the glass, forming volatile silicon fluoride and salts of hydrofluosilicic acid;
but the latter are decomposed by the concentrated sulfuric acid into sulfate,
hydrofluoric acid, and silicon fluoride,

Na,CaSi;O.,+28HF =14H,0 Tf +Na,SiF,+CaSiF,+4SiF, f ,

Soda glass
and

NaSiF;+H.SO, =Na,.SO,+2HF T +Sik, T ;
CaSiF,+H.SO, =CaSO,+2HF f +SiF, ft.

408 REACTIONS OF THE ACID CONSTITUENTS

The silicon fluoride formed by this reaction is a colorless gas with a pene-

trating odor, and is decomposed by water, forming gelatinous silicic acid and
hydrofluoric acid,

(1) SiF,+4H.0 =H,Si0.+4HF,

Silicon fluoride, however, readily combines with hydrofluoric acid to form
hydrofluosilicic ‘acid:

(2) SiF.+2HF =H,[SiF 4,

and the latter compound is not decomposed by water. The whole reaction,
therefore, which takes place between silicon tetrafluoride and water is expressed
by the sum of the two equations, as follews:

38SiF.+ 4H.0 = H,Si0,+ 2H.[SiF’].

If, therefore, a fluoride be heated in a glass test-tube with concentrated
sulfuric acid, and the escaping vapors allowed to act upon water by placing
a moist glass rod in the tube, the water adhering to the rod will become
turbid.

Remark.—Although the above test rarely fails when relatively large
amounts of fluoride are used, it will not be obtained in the case of certain
minerals containing fluorine, such as topaz, tourmaline, etc. The test may
fail, furthermore, if the fluoride is mixed with a large excess of that modifica-
tion of silicic acid which is most readily attacked by hydrofluorie acid. Ac-

cording to Daniel,* this is due to the formation of a stable oxyfluoride, prob-

ably of the formula SiOF>.
The silicon tetrafluoride at first formed combines with the excess of amor-
phous silicic acid present, as follows, .

SiF,+Si0, =2Si0F,,

but this reaction will take place only very slowly, if at all, with quartz powder
or with the silica of a silicate such as glass.

A positive result will be obtained invariably when the tetrafluoride test is
made in a platinum vessel with a relatively large amount of fluoride and
comparatively little amorphous silicic acid or silicate (large amounts of quartz
do not influence the reaction); the test will be negative, on the other hand,
if made in platinum with no silicic acid, or, strange to say, when only quartz
is present with the fluoride. The reason for this different behavior lies 1 in the
difficulty with which quartz is attacked by hydrofluoric acid.

Daniel recommends the following method for performing the test:

Mix the substance to be tested for fluorine with about three times as much
(by volume) ignited quartz powder, place it in a test-tube and stir it into a
thin paste with concentrated sulfuric acid. Close the test-tube with a cork
in which one hole has been bored and an opening cut in the side. Through
the hole in the cork pass a glass rod blackened with asphalt paint, and on
the bottom of the rod suspend a drop of water; push this rod down into the
tube until it is only a distance equal to about 14 times the diameter of the
test-tube from the paste in the bottom. Gently fine the tube and its

*Z. anorg. Chem., 38 (1904, 299).

‘ «A
ee

fh,
i ie

Tae Te

HYDROFLUORIC ACID 409

over a small flame, and if a fluoride is present a white film of H,SiO, will be
formed in the drop of water and will be shown plainly in contrast to the black
rod. In a tube with 1 cm. diameter, fluorine equivalent to 1 mgm. CaF, may
be detected, while with a tube of only 0.5 cm. diameter as little as 0.1 mgm.
CaF, will give the test. When using a tube of small diameter, it is best to add
the sulfuric acid through a small capillary pipette to avoid wetting the sides of

the tube. 7 :

If the substance contains considerable amorphous silica, or when an
oxyfluoride, such as topaz, is present, which is hard to decompose with sulfuric
acid, the test will fail, and it is then necessary to make use of the etching test.

3. The Etching Test.—Place the substance to be tested for fluoride in a
platinum crucible, add some concentrated sulfuric acid, and cover the crucible
with a watch-glass whose convex side has a thin coating of beeswax through
which a few letters have been scratched with a pointed match. On warming
the contents of the crucible, the glass will be etched at the places where the
escaping gas comes in contact with it, if a fluoride was originally present. By
covering the upper concave side of the watch-glass with a little cold water,
the wax coating will not melt during the experiment.

If it is desired to detect the presence of a trace of fluorine, allow the crucible
to stand covered with the watch-glass for twelve hours and then heat for a few
minutes. The presence of only 0.3 mgm. CaF, is sufficient to give this test,
provided a crucible of the right size is used.

If the fluoride contained silica (as in topaz, tourmaline, and other minerals),
the etching test will be negative, for even if the fluorine escapes, it will be in
the form of silicon fluoride, which does not attack glass.

To detect small amounts of fluoride in silicates, it is necessary first to
transform the fluorine into calcium fluoride and to subject the latter com-
pound to the test.

To obtain the fluorine as calcium fluoride, proceed as follows:

Fuse the finely pulverized silicate with six to eight times as much sodium
carbonate in a platinum crucible, and treat the melt with water after it is cold.
A solution is thus obtained in which all of the fluorine is present as sodium
fluoride, together with sodium silicate. Precipitate the silicic acid by adding
considerable ammonium carbonate to the solution, warming it slightly, and
allowing it to stand twelve hours. After filtering off the silica, evaporate
the solution to a small volume, and add a little phenolphthalein, which will
impart a pink color to the solution on account of its being slightly alkaline.
Carefully add hydrochloric acid until the stirred solution becomes colorless,
and heat to boiling, when the color will reappear. Again decolorize the solution
with hydrochloric acid after it has becomes cold, and repeat the process until
the solution becomes only faintly colored on boiling it.

Now add calcium chloride solution and again boil the solution. The
precipitate formed consists of calcium carbonate and calcium fluoride; filter it
off, wash, dry, and ignite it in a platinum crucible. Treat the ash with dilute
acetic acid, evaporate to dryness, triturate with water and filter off the
undissolved calcium fluoride. After drying the precipitate and burning the
filter, it is ready for the etching test.

4. Silver Nitrate causes no precipitation from solutions of soluble
fluorides.

410 REACTIONS OF THE ACID CONSTITUENTS

5. Barium Acetate precipitates barium fluoride soluble in an —

excess of mineral acid and in ammonium salts. Traces of fluorine
present as preservative in foods, liquors, ete., may be detected by
adding a little potassium sulfate (about 0.38 gm.) to the solution,
heating it to boiling and slowly introducing 10 ce. of 10 per cent barium
acetate solution. The precipitate of barium sulfate and fluoride is
then subjected to the etching test.*

6. Calcium Chloride gives a white, slimy precipitate, difficultly
soluble in hydrochloric and nitric acids, but almost entirely insolu-
ble in acetic acid. On account of its slimy consistency, the precipi-
tated calcium fluoride is extremely hard to filter; but by precipitating

can be readily filtered. After igniting and treating with acetic
acid, the precipitate is changed to soluble calcium acetate and in-
soluble calcium fluoride; it is now much denser and can be filtered
readily.

7. Ferric Chloride produces in concentrated solutions of alkali
fluorides a white, crystalline precipitate corresponding to the gen-
eral formula M;[FelF’¢]. These salts, which are analogous to cryo-
lite, Nas[AlF], ave difficultly soluble in water, and their saturated,
aqueous solutions do not give the iron reaction upon the addition of
potassium thiocyanate, except after the addition of acid. These
complex fluorides also are slightly decomposed by ammonia, forming
a basic ferric fluoride.

Methods for Getting Insoluble Fluorides into Solution

(a) Calcium fluoride alone cannot be completely decomposed by fusing
with sodium carbonate. The aqueous solution of the melt always contains
a considerable amount of sodium fluoride, but never the total amount of the
fluorine. If, however, the fluoride is mixed with silica or a silicate, complete
decomposition can be effected by fusing with sodium carbonate. The silica
decomposes calcium fluoride, forming calcium fluorsilicate and calcium silicate,
salts which are decomposed by fusion with sodium carbonate.

On treating the melt with water, sodium fluoride and sodium silicate go
into solution, while the calcium is left behind on the carbonate, and can be
dissolved by treatment with dilute hydrochloric acid.

(b) All fluorides are decomposed by heating with concentrated sulfuric
acid, being changed to sulfates.

* Cf. Buarez, Chem. News, 91, 39; also WoopMAN and Taxpor, J. Am. Chem.
Soc., 27, 1437.

nee oe ee ae

HYDROFLUOSILICIC ACID 411

REACTIONS IN THE DRY WAY

Most fluorides are unchanged by ignition. By heating them
with silica in moist air, they are all more or less completely decom-

posed:
CaF2.+H20+Si02 =CaSi0;+2HF T ;

4HF-+SiO2 =SiF4+2H20.

The acid fluorides give off hydrofluoric acid on ignition, whereby
the glass tube in which they are heated becomes etched.

HYDROFLUOSILICIC ACID, H2SiFs

As we have seen, this acid is formed by the action of silicon fluo-
ride upon water:

38iF4+4H.20 = 2H2[SiF6]-+HuSiOx.

If the silicic acid is filtered off, a strongly-acid solution is obtained
containing hydrofluosilicic acid. By evaporating the solution, the
acid is decomposed into silicon fluoride and hydrofluoric acid,

H[SiFs] =SiFs t +2HF ¢,

so that hydrofluosilicic acid itself is known only in aqueous solution,
although its salts are very stable.

Solubility—Most silicofluorides are soluble in water; the potas-
sium and barium salts form exceptions, being difficultly soluble in
water and insoluble in alcohol.

REACTIONS IN THE WET WAY

A solution of sodium silicofluoride should be used.

1. Dilute Sulfuric Acid causes only a very slight decomposition.

2. Concentrated Sulfuric Acid decomposes all silicofluorides,
evolving silicon fluoride and hydrofluoric acid:

’ NaoSiF. + H2eSO4 = Na2SO1+Sik 4 t +2HF + :

If the reaction is performed in a platinum crucible, the escaping
gas will etch glass, and will cause a drop of water to become turbid.

3. Silver Nitrate produces no precipitation.

4. Barium Chloride gives a crystalline precipitate (1 gm. BalSiFs]
dissolves in about 3750 cc. of water at 17°).

5. Potassium Chloride produces, from solutions which are not
too dilute, a gelatinous precipitate of potassium silicofluoride, which

412 REACTIONS OF THE ACID CONSTITUENTS

is difficultly soluble in water (1 gm. of Ko[Sik’g] dissolves in 835 ce. of
water at 17°) and much more insoluble in an excess of potassium
chloride or in alcohol, but soluble in ammonium chloride.

6. Ammonia decomposes all soluble silicofluorides, with separa-
tion of silicic acid: =

NagSil’s +4NH40H = 2NaF+4NH4F+H,Si04.
7. Potassium and Sodium Hydroxides react in the same way as
ammonia, but the silicic acid remains in solution as alkali silicate. —
REACTIONS IN THE DRY WAY

All silicofluorides are decomposed on being heated into fluoride
of the metal and silicon fluoride:

K2SiF6=2KF+SiF4 T .

The escaping gas renders a drop of water turbid, and the residue
gives all the reactions of a fluoride 7

GROUP VII

NON-VOLATILE ACIDS WHICH FORM SOLUBLE SALTS
WITH THE ALKALIES

SILICIC ACID, H4SiO, AND H2SiO3

Occurrence.—The above acids, from which very stable salts are
derived, are not known in the free state, as is the case with carbonic
and sulfurous acids; although there are indeed amorphous, natural
minerals consisting of hydrated silica with varying amounts of
water: water opal with about 36 per cent water, ordinary opal with
from 3 to 13 per cent water, and hyalite with about 3 per cent water:
but none of these substances represents a compound of constant
composition.

The anhydride SiOz occurs in rhombohedral crystals as quartz,
whose prismatic faces are almost always striated horizontally; and
as tridymite, also crystallizing in the hexagonal system. The amor-
phous silicic acid is often found mixed with the crystallized anhy-
dride as flint, agate, chalcedony, jasper, etc. Silicic acid is, however,
most frequently found in the form of its salts, the silicates.

Preparation and Properties——Silicic acid can be very readily ob-
tained pure by the hydrolysis of its fluoride,

38iF4+4H20 = 2H2SiF's6+H4SiOng,
or by the decomposition of alkali silicates (water-glass) with acids:
NazSi03+2HCl=2NaCl+H2SiOsz.

The silicic acid thus obtained forms an amorphous, gelatinous mass,
appreciably soluble in water and acids, and readily soluble in even dilute
solutions of caustic alkalies or alkaline carbonates. Thus freshly precipitated
silicic acid will be readily and completely dissolved by a short digestion with 5
per cent (or even 1 per cent) sodium carbonate solution on the water-bath.
On being dried, silicic acid gradually loses water, and at a gentle red heat is
changed into the form of its anhydride. According to the extent to which the
dehydration has gone, the solubility of the silicic acid diminishes both in acids
and in alkalies. .

1. Air-dried silicic acid, with 16.65 per cent of water, corresponding to the
formula 3Si0.-2H;0, is perceptibly soluble in acids, and completely dissolved
by digestion for one-quarter to one-half an hour with 1 per cent soda solu-
tion on the water-bath.

413

414 REACTIONS OF THE ACID CONSTITUENTS

2. Silicie acid dried at 100° with 13.60 per cent of water, corresponding
to the formula 2Si0,-H,O, is practically insuluble in acids, but can be dissolved
by digesting for one-quarter hour with 1 per cent sodium carbonate solution
upon the water-bath, or more readily by boiling.

3. Silicic acid dried at 200°, with 5.66 per cent of water, corresponding
to the formula 5SiO.-H.O, and the acid dried at 300°, with 3.40 per cent of
water, corresponding to the formula 9S8i0.-H.O, dissolve slowly by digestion
with i per cent sodium carbonate solution on the water-bath.

4. The anhydride obtained by gentle ignition to a faint-red heat is only
partly dissolved by 1 per cent or by 5 per cent sodium carbonate after half
an hour’s digestion on the water-bath; but is dissolved after boiling for two
hours with the sodium carbonate solution.

5. The strongly ignited anhydride is dissolved slowly by 5 per cent sodium
carbonate solution after repeated boiling for a long time, but is readily dis-
solved by boiling with concentrated caustic soda or potash.

6. The native anhydride, quartz, after being powdered in an agate mortar,
is practically insoluble in 5 per cent sodium carbonate solution, and very diffi-
cultly soluble in boiling caustic alkali. If it is in the form of an extremely
fine powder, it can be dissolved by boiling with 5 per cent sodium carbonate
solution (Lunge and Millberg).

It follows from the above that the solubility of silicic acid (and of its
anhydride) in alkali carbonates depends largely upon the fineness of the material.

Silicie acid, as well as its anhydride, is soluble in aqueous hydrofluoric
acid, forming hydrofluosilicic acid:

Si0.++6HF =2H.0+H.SiF,.

By evaporating this solution hydrofluoric acid is evolved; and silicon
fluoride, with small amounts of silicic acid, is left behind. In order, then,
to volatilize silicic acid completely by means of hydrofluoric acid, the hydro-
lytic action of water must be prevented, which is effected by the addition of
a little concentrated sulfuric acid. The procedure is as follows:

Moisten the substance in a platinum crucible with a very little water, add
not more than 3 ec. of concentrated sulfuric acid, and then the hydrofluoric
acid. Evaporate the mixture on the water-bath, or suspend the crucible in a
larger crucible, and heat the latter until the hydrofluoric acid is expelled, cool,
add another portion of hydrofluoric acid and again evaporate. If a very
large quantity of silicic acid is present it may be necessary to treat with
hydrofluoric acid a third time. It is better to proceed in this way than to
add a large quantity of hydrofluoric acid at one time. Finally drive off the
sulfuric acid by heating directly over a small flame.

The salts of silicic acid, the silicates, are exceedingly numerous, and are
usually very stable. Many of them are so stable that they are not attacked
by concentrated acids, while others are easily decomposed thereby.

The different silicates are classified according to their solubility into

A. Water-soluble silicates.

B. Water-insoluble silicates, which are again divided into

(a) Silicates decomposable by acids;
(b) Silicates undecomposable by acids

SILICIC ACID 415

A. Water-soluble Silicates

The silicates which are soluble in water, or ‘‘ water-glasses,” are obtained
; £ ’

by fusing silica or a silicate with caustic alkali or alkali carbonate:
Si0.+ NaCO; =Na.Si0;+CO, i} e

1. Behavior Toward Acids.—The aqueous solution of an alkali
silicate reacts strongly alkaline, showing that the salt is hydrolyzed
to a marked degree:

Si03° +2H20 — 20H + H2Si03.

The silicic acid set free by the hydrolysis is present as hydrosole
in the solution. By the addition of acid the alkali hydroxide is con-
verted into salt and a part of the silicic acid is coagulated, provided-the
solution is not too dilute.

The precipitation is by no means quantitative; a considerable quantity
of silicic acid remains in solution and, in fact, under some conditions all of it
may remain dissolved in the dilute acid. If, namely, a 10 per cent water-glass
solution is poured quickly into hydrochloric acid of specific gravity 1.1 to
1.3, there*is no precipitation. After standing some time, however, the entire
contents of the beaker are changed to a jelly. From 1 per cent solutions
treated similarly with acid, no precipitate appears even after standing a year.

The silicic acid which is precipitated upon the addition of acid is, there-
fore, considerably soluble in dilute acids. In order to separate the silicic
acid completely from a solution of water-glass, the hydrated acid must be
changed into the less hydrated acid, 28i0.-H.O, by heating at 100° C. (ef.
p. 414). For this purpose acidify the water-glass solution with hydrochloric
acid (or nitric or sulfuric acid) and evaporate on the water-bath to complete
dryness (the mass must no longer smell of acid). Moisten the dry residue
with strong acid, warm slightly, dilute and filter off the silicic acid. The small
amount of silica remaining in solution can be removed almost entirely by a
second evaporation of the filtrate.

2. Behavior Toward Ammonium Salts.—If a solution of water-
glass is treated with an ammonium salt, the silicic acid will, for the

‘most part, be precipitated as hydroxide; the precipitation is not quite

quantitative, but more complete than is obtained by the addition
of cold dilute acid:

Si03~+2NH4* oe 2NH3+H2SiO3;
NH3+H20 @ NH40OH.

The hydroxyl ions have a marked solvent action upon the silicic
acid. For this reason the precipitation is more complete with an
ammonium salt of a strong acid than with that of a weak acid, which

416 REACTIONS OF THE ACID CONSTITUENTS

is already hydrolized to a considerable extent. Boiling off the
ammonia helps to make the reaction complete. The use of ammonium
carbonate, though less satisfactory than ammonium chloride, is
necessary when it is desired to test the solution for chloride.

Silicic acid is more completely precipitated by zinc-ammonia
hydroxide than by ammonium carbonate,

Si03° +[Zn(NH3)6](OH)2=20H +6NH3+ZnSiOsz,

because the zinc silicate formed by the reaction is much more difficultly
soluble in dilute alkaline solution than is the free silicic acid.

The separation of silicic acid from a solution of water-glass by means of
ammonium carbonate may be illustrated -by a common case. Many rocks
(particularly the zircon-syenite of Norway and Greenland, many granites and
basalts) contain small amounts of sodalite, NaCl-3NaAISi0,, a chloride silicate
of the leucite group. In order to detect the chlorine in such a rock, the fol
lowing process may be used: Fuse the finely powdered silicate with six
times as much sodium carbonate in a platinum crucible, extract the product of
the fusion with cold water and filter. The filtrate contains all of the chlorine
as sodium chloride in the presence of sodium silicate. Treat the solution
with ammonium carbonate, warm gently, allow to stand twelve hours, and
then filter off the precipitated silicic acid. In order to separate the rest of
the silicic acid, add a little zinc-ammonia hydroxide and boil the solution
until it no longer smells of ammonia. Filter off the precipitated zine silicate
and zinc oxide, acidify the filtrate with nitric acid and test for chlorine with
silver nitrate.

To prepare the zinc ammonia hydroxide dissolve pure zine in nitric acid,
treat the solution with potassium hydroxide solution until it is neutral, and
dissolve the filtered and washed zinc hydroxide in 6-normal ammonium
hydroxide.

B. Silicates Insoluble in Water
(x) Decomposable by Acids

A large number of native silicates are decomposed by evapora-
tion with hydrochloric acid, the silica being deposited sometime.

in the form of a jelly and sometimes in the form of a powdery mass. _

All zeolites, and a number of artificial silicates (such as Portland
and Roman cements) belong to this class of silicates.

To remove all of the silicic acid from these silicates, treat the finely powdered
mineral with dilute hydrochloric acid, evaporate to dryness on the water-
bath, moisten the mass with concentrated hydrochloric acid to convert any
oxides or basic salts of iron, aluminium, magnesium, etc., into soluble chlorides,
heat gently, dilute with hot water, boil and filter. The silicic acid is left on
the filter, and the filtrate contains the metals as chlorides (ef. p. 414).

The purity of the residual silicic acid must always be tested. For this
purpose, place the well-washed precipitate together with the filter-paper, in

SILICIC ACID AIT

a clean platinum crucible, held in an inclined position on a triangle, and care-
fully burn the filter-paper. ‘Treat the residue with water, concentrated
sulfuric acid and hydrofluoric acid as directed on p. 414 and finally remove
the excess of sulfuric acid by cautious heating over the free flame. If the
silicic acid were pure, nothing should remain after the evaporation of the
sulfuric acid. Almost always a small residue of aluminium and ferric oxides
remains, which in most cases can be neglected. If considerable residue is left,
it should always be tested for titanic acid, barium sulfate and possibly tin
dioxide.

(8) Silicates Undecomposable by Acids

Most silicates, the feldspars, micas, artificial glasses, porcelain,
etc., belong to this class. In order to remove the silicic acid from
such substances, they must be

1. Fused with an alkali carbonate.
2. Fused with lead oxide or boron trioxide, or
3. Heated with sulfuric and hydrofluoric acids.

Silicic acid is not soluble in water except to form colloidal solutions, and
yields scarcely any hydrogen ions. In this respect it is a very weak acid but,
as it is practically non-volatile, it is capable of expelling the acid from the
salts of strong acids provided the base itself is not volatile at the temperature
at which the salt is decomposed. The silicates, therefore, are very stable
compounds particularly toward heat. ‘The natural silicates are partly derived
from ortho silicic acid, H.SiO., and partly from meta-silicic acid, H.SiO;, but
like other polybasic acids, silicic acid also forms salts which are derived from
polysilicic acids such as H,48i;03, and Hi2SisO1. The solubility of the silicate
depends upon two factors—the solubility of the oxide of the base and the
proportion of silicic acid which it contains. As a general rule, the salts of
ortho- and meta-silicic acids are more soluble than those of the polysilicic
acids. Thus sodium and potassium ortho- and meta-silicates are soluble in
water whereas a polysilicate may contain alkali as its principal base and yet
be undecomposable by concentrated hydrochloric acid. Such silicates may,
however, be decomposed by hydrofluoric acid, which causes volatilization of
the silica as silicon tetrafluoride, by melting with a solid acid such as boric
acid, or by treating with mineral acid in a sealed tube.

Salts of bases of which the ignited oxides are very insoluble, such as Al,O,,
sometimes form insoluble silicates of the ortho and meta types.

- The effect of fusing a silicate with an alkali carbonate, or with a fusible
oxide of some metal such as lead, is to increase the proportion of base in the
silicate molecule. When the proportion of base-is increased, the solubility of
the silicate is also increased provided the base is itself readily soluble in acid.
It is not at all necessary, therefore, to get all the silicic acid in the form of
sodium silicate or of lead silicate, by fusing with sodium carbonate or with
lead oxide, but it is sufficient if the silicate is converted into a silicate which
is decomposable by acid. For this reason the fusion with sodium carbonate
or with lead oxide is often said to open up the silicate. It converts the silicate
into the ortho or meta type and makes the silicate decomposable by acid.
Thus after fusing with sodium carbonate, for example, it will be found that

418 REACTIONS OF THE ACID CONSTITUENTS

part of the sodium is converted into water-soluble silicic acid and part of it
into a double silicate which is decomposable by acid. <A part of the sodium
and a part of the silicate can be dissolved out of the fused mass by treatment
with hot water. |

1. Fusion with an Alkali Carbonate—This method is commonly used when
it is desired to detect the presence of silicic acid and of all the bases except
the alkalies.

Mix the finely powdered substance with 4—6 times as much calcined sodium
carbonate (or a mixture of equal parts of sodium and potassium carbonates,
which melts lower than sodium carbonate alone), and fuse the mixture in a
platinum crucible, heating carefully at first to avoid spattering from too violent
evolution of carbon dioxide. Gradually increase the temperature until the full
heat of the burner is reached, and continue fusing until the molten mass is
quiet, and then heat for about a quarter of an hour over the blast lamp.
Make a spiral by winding some platinum wire around a stirring rod, and insert
the spiral in the melt. Cool the crucible quickly by directing a blast of cold
air against its sides, and while the contents of the crucible are still warm, but
not hot enough to spatter badly on the water, cover with a little water from
the wash bottle. After a few minutes the fusion can usually be withdrawn
with the aid of the platinum spiral. Treat the product of the fusion as described
under (a).

2Na.Si0;+CaCO;+2Al0(ONa)+14HCl =

=6NaCl+CaCl,+2AICl;+2H.S8i0;+CO, T +5H.0.

During the evaporation to dryness salts like aluminum chloride (ferric chloride,
ete.) are subject to hydrolysis and are converted to some extent into oxide or
basic salt, insoluble in water. Therefore, in order to separate the silicic acid
from the salts, it is first necessary to convert such oxides or basic salts back
into chlorides. This is accomplished by moistening the dry residue with con-
centrated hydrochloric acid. After warming the acid with the residue for
about ten minutes, dilute with hot water, boil and filter off the silicic acid using
an ashless filter.

To identify the silicic acid, place the well-washed precipitate, together with
the filter, in a weighed platinum crucible, dry cavefully by a low flame placed
in front of the crucible, and then ignite at as low a temperature as possible,
with the flame now at the base of the crucible, until the carbon of the filter
is all consumed. Then, for the first time, ignite strongly, eool somewhat,
place in a desiccator and weigh when perfectly cool. Treat with hydrofluoric
and sulfuric acids as described on p. 414, ignite and weigh after proper cooling.
A difference in the weights before and after the treatment with these acids
shows not only the presence of silica, but also the quantity of it. This quan-
titative method is necessary for the detection of small quantities of silicic acid.

To identify the silicic acid qualitatively, Daniel’s tetrafluoride test is satis-
factory.*

Daniel’s Tetrafluoride Test

Ignite the well-washed precipitate, as described above, in a platinum cru-
cible, then triturate in a mortar with three times as much potassium-sodium
carbonate, and fuse the mixture in the crucible. After cooling the melt,

* Z. anorg. Chem., 38, 299 (1904).

SILICIC ACID 419

soften it by heating with a little water and treat with dilute sulfuric acid to
decompose the excess of carbonate as well as the salt of silicic acid formed during
the fusion. Heat the mixture in the crucible, by placing the latter upon a
piece of asbestos board, and evaporate nearly to dryness, or until a thick
jelly of silicic acid remains. After cooling, add three times as much fluorspar
as there was original precipitate, a little magnesite and enough concentrated
sulfuric acid to make a thin paste. After mixing the contents of the crucible
with the aid of a stout platinum wire, place a drop of water on the inside of
a crucible cover, which is partly painted with asphaltum, place the cover on
the crucible and heat the contents of the latter gently. From time to time,
raise the cover to see whether the water has become turbid. It frequently
happens that the water becomes turbid, and then, provided a large excess of
hydrofluoric acid is present, the turbidity disappears. For this reason the cover
must be inspected frequently in order not to miss any temporary turbidity.

The tetrafluoride test for silicic acid is very sensitive if the reaction is
carried out in a very small platinum crucible. If such a crucible is not at
hand, with a capacity of say 0.5 to 1 cc., it is better to test by the quanti-

tative method when less than 0.01 gm. of silicic acid is present.
| 2. Fusion with Lead Oxide or Boron Trioxide——These methods are very
rarely used in qualitative analysis, so that it will not be necessary to describe
them here. They play a more important part in quantitative analysis and
will be described, therefore, in the second volume of this book.

3. Decomposition by Hydrofluoric Acid.—This method is used principally
when a silicate is to be examined for alkalies, titanic acid or barium. Treat
the finely powdered silicate in a platinum dish with about 2 ce. of pure sulfuric
acid (1 vol. concentrated acid and 2 vols. of water) and about 5 cc. of freshly
distilled hydrofluoric acid and evaporate the mixture on the water-bath, stirring
the mass from time to time with a thick platinum wire until it no longer smells
of hydrofluoric acid. Add 5 cc. more of hydrofluoric acid and again evaporate,
finally heating the dishf very carefully over the free flame, under a good hood,
until the greater part of the sulfuric acid is expelled. The mass should not
be ignited strongly, for a part of the sulfate may then be changed to an oxide
insoluble in water. The sulfates of iron and aluminium, for example, are
decomposed on ignition. After cooling, treat the mass with water, and usually
everything will gradually go into solution. If a residue remains, test it for
barium sulfate and titanic acid. . The solution can be used for the alkali tests,
or for the tests for the other metals, if it is desired.

REACTIONS IN THE DRY WAY

If silicic acid or a silicate is heated in the salt of phosphorus bead,
the metallic oxide will dissolve, while the silicic acid itself will be left
as a white gelatinous mass, suspended in the bead (skeleton bead).
This reaction, however, is, not infallible for certain silicates of the
zeolite group dissolve in the bead without the formation of the skeleton.

SILICON, Si. At. Wt. 28.3

Silicon exists in two modifications, one of which is crystalline, while the
other is amorphous. Amorphous silicon is a dark-brown powder, which can
be oxidized by heating in the air while the crystalline modification remains
unchanged on ignition in pure air or in oxygen, but if the air contains carbon
dioxide, it is oxidized to silicon dioxide with deposition of carbon:

CO.+8i =810.+C.

Crystallized silicon is not attacked by any acid, but is readily dissolved by
boiling with concentrated caustic alkali with evolution of hydrogen:

Si+20H-+H,0 <> Si0;-+ 2H, tT .

Silicon unites with many metals, forming silicides. The silicides of the
light metals, magnesium, calcium, etc., are decomposed by dilute hydrochloric
acid with the formation of spontaneously combustible silicon hydride:

Mg.Si+4H+ > 2Mg+++H.si f.

The hydride of silicon is not spontaneously combustible when pure, only
when it is contaminated with hydrogen, as is invariably the case.

In order to detect the presence of silicon in such a compound, treat it with
nitric acid, which oxidizes the greater part of the silicon to silicic acid.

Detection of Silicon in Iron and Steel

If it is a question of detecting the presence of silicon in the different kinds
of irons (steel, cast iron, etc.) take a large amount of material, for the amount
of iron silicide present is usually very small. Place 5 or 10 gms. of the material
(best in the form of borings) in a large beaker and treat with 60 cc. of 6-normal
nitric acid. A violent reaction at once takes place with evolution of brown
nitrous fumes. As soon as this action lessens, heat the solution to boiling,
and continue heating until no more brown fumes are given off. Then pour
the solution into a 200-cc. casserole and evaporate as far as possible upon the
water-bath. Heat the residue carefully over a free flame until it is perfectly
dry and then ignite the mass until the nitrate is completely changed to oxide,
when no more brown fumes will be evolved. After cooling, dissolve the mass
in about 50 cc. of concentrated hydrochloric acid, heat with constant stirring
almost to boiling, evaporate nearly to dryness, take up in water, filter, and test
the residue for silicic acid, by seeing whether it is volatile with sulfuric and
hydrofluoric acids.

In the case of cast iron, the silicic acid obtained is usually considerably
contaminated with graphite, which can be removed by long ignition in a
platinum crucible before treating with hydrofluoric and sulfuric acids.

420

SILICON 421

Detection of Silicon in Carborundum and Metal Silicides

Other silicides, such as carborundum, SiC, are not decomposed by nitric
acid; they can be fused with caustic alkali in a silver crucible,

SiC+4KOH+2H,0 =K.Si0;+ K,CO;+4H, tT ;

and on acidifying the melt, the silicic acid separates out.

Carborundum in the form of a fine powder is also easily decomposed by
fusing with potassium carbonate. On removing the cover of the platinum
crucible the blue flame of burning carbon monoxide is seen:

3K,CO;+8iC =K,.Si0,+2K,0+4C0 f.

The method of fusing silicides with caustic alkali is often used for getting
metallic silicides into solution. Many copper silicon alloys are scarcely attatked
_ by even aqua regia. If, however, they are fused with caustic alkali in a sil-
ver crucible, potassium silicate, metallic copper, and hydrogen are formed:

SiCu.+20H-+H,0 =Si0;-+Cu,.+2H: f .

By treating the melt with water, the soluble potassium silicate may be separated
from the copper. .

PART IV. SYSTEMATIC ANALYSIS

THE purpose of a qualitative analysis is not simply to find out
what elements are contained in a given substance, but the aim should.
also be to get a good idea of the relative amounts that are present.
Manganese chloride, for example, is made from pyrolusite, and
almost always contains traces of calcium, magnesium, nickel, cobalt,
and iron. If the analyst should report that ‘the analyzed sub-
‘stance consists of chlorides of calcium, magnesium, nickel, cobalt,
iron, and manganese,” it is evident that one would get but a poor
idea of the nature of the substance. .The report should read: ‘‘ The
substance examined was manganese chloride, and contained traces of
calcium, magnesium, etc., as impurities.”

In order to be able to estimate the relative amounts of the differ-
ent components of a substance, it is necessary to start with a known
amount (usually } to 1 gm.) and compare the size of the precipitates
produced. It will be impossible for the beginner to estimate the
amount of a precipitate obtained, if he has studied the reactions of
the elements with unknown amounts of the different substances.
If, however, he has learned to work with a known amount of material,
he will soon be able to judge from the size of a precipitate the
amount of element to which it corresponds.

It is a good plan first to work through the analysis of each group
with a known solution containing 10 mgms. of each element and then
it is comparatively easy to determine approximately how much of each
element is present by the test obtained in the analysis of any unknown.
Thus, starting with 1 gm. of the original substance, it is often con-
venient to designate as present in small quantity when apparently
less than 10 mgms. is found, as present in medium quantity when from
10 to 50 mgms. is found, and as present in large quantity when distinctly
more than 50 mgms. is present. Experiments with large classes of
students have shown that such judgments are correct in nearly nine
cases out of ten. It should be borne in mind, moreover, that for
estimating small quantities of substances, qualitative tests are more
accurate than any method of quantitative analysis. Thus all col-
orimetric methods of quantitative analysis are really based on
qualitative tests. The comparison of a test with one obtained using
a known quantity of substance often gives a more exact determination

422

PRELIMINARY, EXAMINATION 423

of the quantity present than a method involving weighing or titra-
tion. This is because it is easy to prepare a solution containing say
1 mgm. per cubic centimeter with an error of less than 5 parts per
thousand, by dissolving 1000 times as much in a liter and thoroughly
shaking; but it is more difficult to determine with equal accuracy
the presence of only 1 mgm. of substance. By the process of diluting
and taking an aliquot part, it is possible to prepare a solution contain-
ing a very small known quantity of any soluble substance. For
convenience, it is well to have solutions at hand containing exactly
10 mgms. per cubic centimeter of each constituent. By taking three
small drops of such a solution, approximately 1 mgm. of the constitu-
ent can be obtained.

Every analysis should be divided into three parts:

I. The preliminary examination.

II. The examination for the metals (cations).
III. The examination for the negative elements (anions).
The substance analyzed may be

A. Solid and non-metallic.

B. A metal or an alloy.

C. A solution (liquid).

D. A gas.

The whole amount of the substance at hand should never be used
for the first analysis, but a portion should always be reserved for
unforeseen accidents. The portion taken for analysis should be divided
into two parts after the preliminary examination, the first part being
used for the tests for the electro-positive and the other part for the
tests for the electro-negative elements.

Before beginning an analysis, the substance should be carefully
examined with the naked eye and with the microscope, and the results
noted. Oftentimes the odor, color, and crystalline form suffice to
give important clues as to the nature of the substance.

A. THE SUBSTANCE IS SOLID AND NON-METALLIC *
I. PRELIMINARY, EXAMINATION

This should never be omitted, for it often shows how the subse-
quent analysis may be considerably shortened, and in some cases
makes the further examination unnecessary. It consists only of
making the following few simple tests:

1. Heating in the Closed Tube.—By a closed tube is understood
a small glass tube about 10 cm. long and 0.5 cm. in diameter sealed at

*See p. 452 for B and C.

424 SYSTEMATIC ANALYSIS

one end. Place a little of the substance in the tube so that none of
it remains adhering to the sides, hold the tube in a nearly horizontal _
position and cautiously heat in the flame, noting carefully whether
any change takes place.

The Substance is Volatile

(a) The substance sublimes completely without any deposition of
water; it contains no non-volatile substance.

The sublimate is white. The halogen compounds with ammonia,
mercurous chloride and bromide, mercuric aminochloride, arsenic
trioxide and arsenic pentoxide may be present.

Arsenic pentoxide melts before being changed into the trioxide.

The sublimate is colored—
Gray: all oxygen compounds of mercury, cyanide of mercury,
free iodine, and arsenic.

Mercuric cyanide leaves a brown mass, paracyanide, which only disappears
after long-continued heating.

Yellow: arsenic sulfide, sulfur, mercuric iodide.

Mercuric iodide becomes red immediately on being rubbed with a glass
rod.

Grayish black: mercuric sulfide.

(b) The substance is completely volatile, with separation of water
and gaseous products: most ammonium compounds (with the excep-
tion of those of the halogens) and free oxalic acid.

By very cautious heating, oxalic acid may be sublimed; it usually de-
composes, however, into water, carbon monoxide, and carbon dioxide.

The Substance is only Partly Volatile

In this case gases and vapors may be evolved:

Oxygen from peroxides, nitrates, chlorates, iodates, ete.

Carbon dioxide from carbonates and organic substances; in the
latter case it is usually accompanied with the separation of carbon
and evolution of empyreumatic, combustible vapors.

Chlorine from chlorides of platinum, gold, copper, iron, ete.

Iodine from iodides, in the presence of oxidizing substances.

Sulfur from many sulfides and thiosulfates.

Arsenic from arsenites and arseniates, in the presence of carbon
or organic substances.

PRELIMINARY EXAMINATION 425

Arsenites are reduced without the aid of charcoal:
-10K;As0; =6K3As0.+-6K,0+Asi.

Water from substances containing water of crystallization, trom
acid salts, organic substances, or from the phosphate, borate, chromate,
vanadate, and tungstate of ammonium.

The water given off condenses in the cooler part of the tube and
should be tested with litmus-paper. If it reacts alkaline, it comes
from ammonium compounds; if acid, it results from easily decom-
posable salts of the stronger acids.

-Many fluorides when heated with water give off hydrofluoric
acid, which etches the glass.

If a sublimate is formed, make the following experiment:

Mix a little of the substance with three times as much calcined
sodium carbonate and heat in the closed tube. If ammonium salts are
present, the smell of ammonia can be detected. Mercury compounds
give a deposit of gray metal (cf. p. 203); arsenic and its oxygen com-
pounds also usually yield the gray metal (but no globules), accom-
panied by a garlic odor.

The oxygen compounds of arsenic do not give the metal when heated with
pure sodium carbonate: Commercial sodium carbonate, however, is usually
contaminated with enough paper fibers to cause the reduction.

2. Test the Substance in the Bead. Make a borax or sodium
phosphate bead in the loop of a very thin platinum wire (as described
on p. 64), introduce it with a little of the substance into the oxidizing
flame, observe the color of the bead both when it is hot and when it is
cold, and then heat it in the reducing flame. Borax is usually used for
this experiment, except when it is desired to test for silicic or titanic
acids, or when the substance is white, in which case salt of phosphorus
is used. Only colored oxides are capable of coloring the borax bead.

Some oxides are reduced to metal, so that the bead appears gray in the
reducing flame (see following table). CuSO, is white when anhydrous, but
becomes blue immediately on the addition of water.

The following substances impart a characteristic color to the bead:
iron, manganese, nickel, cobalt, chromium, uranium, copper (didy-
mium, cerium, vanadium, titanium, and tungsten).

Since the coloration varies with the temperature and with the
amount of substance used, the results to be expected, with the neces-
sary conditions, are summarized in the table on p. 426. The following
abbreviations will be used: h=hot; c=cold; h—c=hot and cold;
s.s.=slightly saturated; sat.=saturated. |

426

SYSTEMATIC ANALYSIS

Wito Borax.

Wits SALT or PHOSPHORUS. —

Color of
the Bead. In the Oxidizing In the Reducing In the Oxidizing In the Reducing
Flame. Flame. Flame. Flame.
SiO. (without |SiO. (without | SiO. (usually | SiO. (usually
skeleton), al-| skeleton), al-| with skeleton),| with skeleton),
kaline earths,| kaline earths] alkalineearths| alkalineearths
Colorless Hg, Pb, Bi,| and earths,}| and earths | and earths,
Sb, Cd, Zn,| Mn, Di, Ce, (sat. = tur-| Mn, Di, Ce,
Sn, Ti Cu (s.s.) bid) Cu (s.s.)
W, Mo, Fe :
(s.s.—c) W, Ti
G Ag, Pb, Bi, Sb, Ag, Pd, Bi, Sb,
sand Cd, Zn, Ni Cd, Zn, Ni
: Fe (s.s.—h), Ag Fe (s.s.—h), Ag
Yellow (h), Ce (h), Uj ,_.. (h), Fe (sat.—
(or th), Vv ogre m c), Ce (h), V|Fe (h), Ti (h)
brown) sat.), Ni (c) ery (h), U (h), Ni
(brown) (c) (brown)
| Cr (c), Cu (h)
Fe (h—c), Ur, : ;1Cr (c),. U (e)
Green Cr (c), Cu (h) Cr. V (h) ai (h), U Vv (ae Mo fey
(c—sat.)
Co (h—e), Co (h—c),
Blue a Ry Co (h—c) Cu [Co the), W (6)
Mn (h—ce), Di :
Violet || (h—c),.and Ni any at Dl oe
(with cobalt) ¢)
Cu (sat.),
opaque; when Cu as in the
very slightly borax bead;
Red Fe (h—sat.) saturated and}| Fe (h—-sat.) Ti and W in
. Ce (h) with a trace of Ce (h) the presence.

Sn, ruby red
and transpar-
ent. :

of iron=blood
red

3. Heat a Little of the Substance upon Charcoal before the Blow-
pipe; if deflagration takes place a nitrate, nitrite, chlorate, iodate,
etc., may be present.

4. Heat the Substance with Soda upon Charcoal before the Blow-
pipe.—Mix as much of the substance as can be taken up on the end
of a knife-blade with twice as much sodium carbonate (as described
on p. 68), place it in a cavity on a piece of charcoal and heat in the
reducing flame of the blowpipe.

PRELIMINARY EXAMINATION 427

There is obtained:

‘As malleable button: Au, Ag, Sn, Cu, which
can be pressed flat in an agate mortar.

As gray metallic particles: Pt, Fe, Ni, and Co.
Pt may be pressed flat in an agate mortar;
Fe, Ni, and Co are magnetic and are attracted
| by a magnet (cf. p. 66).

(As a brittle metallic button: Sb (white in-
crustation), Bi (yellow incrustation). The
button may be reduced to a powder by
grinding in an agate mortar.

As a malleable button: Pb (yellow incrus-
| tation).

White, yellow when hot: Zn.
(c) Incrustation without metal. ; Brown: Cd.
White: As (garlic odor).

(a) Metal without incrustation.

(6) Metal with incrustation... . ;

(d) White, infusible, strongly

luminous mass........... Ca, Sr, Mg, Al, and rare earths.

(e) Sulfur compounds are reduced to sulfides. If the melt is placed on
a bright silver coin and moistened with water, the silver is blackened (Hepar
reaction).

5. Test the Substance to See whether it Imparts Any Color to the
Non-luminous Flame.—Introduce a little of the substance on a plati-
num wire into the base of the flame (cf. p. 62), and then into the
fusing zone. Afterwards moisten it with dilute hydrochloric acid
and repeat the experiment. The following indications may be ob-
tained:

Sodium gives a yellow monochromatic flame; a piece of sealing-
wax or a crystal of potassium dichromate appears yellow when illumi-
nated by this flame.

Potassium (cesium and rubidium) gives a violet flame which is
completely obliterated by the sodium flame. If the flame is observed
through cobalt glass, the sodium flame disappears and the potassium
flame appears pink.

Lithium gives a carmine-red flame (or a red line in the spectro-
scope).

Strontium also gives a carmine-red flame (which the spectroscope
shows to consist of severa! lines in the orange, and a bright line in the
blue).

Calcium gives a brick-red flame (in the spectroscope an orange
and a green line are seen, both about an equal distance away from
the sodium line).

428 SYSTEMATIC ANALYSIS

Barium gives a greenish-yellow flame.

In the case of barium sulfate the green flame is either only indis-

tinctly visible or-not at all. In order to detect barium in this case,
heat a small portion of the substance in the upper reducing flame;
after cooling moisten it with hydrochloric acid (odor of hydrogen
sulfide) and again heat, when the barium flame can be easily seen.

Thallium gives an emerald-green flame.

If a green flame is obtained, test another portion of the substance
for boric acid, by treating with concentrated sulfuric acid and bringing
near the flame. A green color indicates the presence of boric acid,
but if copper is present this test is not reliable.

By heating the solid‘substance with potassium ethyl sulfate in a test-tube,
boric acid is converted into B(OC:H;)s, which is volatile and burns with a
green flame. Copper chloride does not interfere with this test.

Lead, Arsenic, Antimony color the flame light blue, and copper
compounds color the flame either green or blue.

Preliminary Examination for the Electro-negative Elements
(Anions)

1. Dilute Sulfuric Acid (2-normal).—Treat about a gram of the
substance in a small test-tube with dilute sulfuric acid, and note
whether a reaction takes place in the cold or not (evolution of a gas).

The following gases can be recognized:

HCN from cyanides (odor) ;

Mercuric cyanide does not liberate HCN in this test.

H2S from soluble sulfides (odor, and blackening of lead acetate
paper) ;

NOz from nitrites (brown fumes) ;

SO2 without separation of sulfur from sulfites (odor of burning
sulfur) ;

SOz accompanied by separation of sulfur from thiosulfates; the
-deposited sulfur is yellow, particularly after warming; .

COz from carbonates or cyanates (barium hydroxide solution is
rendered turbid). i

By boiling with dilute sulfuric acid, soluble ferro- and ferri-
cyanides are decomposed and evolve hydrocyanic acid; acetates set
free acetic acid; hypochlorites evolve chlorine (which also takes place
in the cold); while the peroxides of the alkalies and alkaline earths
are decomposed with evolution of oxygen.

Alkali peroxides also evolve oxygen when treated with water. Cf. p. 83.

PRELIMINARY EXAMINATION 429

2. Concentrated Sulfuric Acid.—If the substance does not react
with dilute sulfuric acid add 3 or 4 ce. of concentrated sulfuric acid
and heat. If the substance reacted with dilute sulfuric acid, it will
react violently with concentrated sulfuric acid and the gas will come
off so quickly that it will carry small particles of the sulfuric acid with
it, which makes the gas appear to have a penetrating odor and may
lead to a mistaken conclusion, especially as it will also cause barium
hydroxide solution to become turbid.

In such a case, add dilute sulfuric acid drop by drop to a new
portion of the substance until no further action takes place, then
add 5 ce. of concentrated sulfuric acid and heat the mixture.

Gases and vapors may be evolved, which are

(a) Colorless

HCl from chlorides, fuming in the air, with penetrating odor.
The fumes do not cause a turbidity with water.

AgCl and HgCl, evolve HCl very slowly; the same is true of Hg,Cl., and
in this case SQ, is also set free. Cf. p. 288.

SiF4 from fluorides, fuming in the air, with a penetrating odor,
and causing a turbidity on coming in contact with water.

SiF, is formed on account of the experiment being performed in glass. In
platinum and in the absence of silica, HF would be evolved, which does not
render water turbid.

SOz, without separation of sulfur. If there was no evolution of
sulfur dioxide on treatment of the substance with dilute sulfuric acid,
the sulfur dioxide which now escapes must come from the sulfuric
acid itself; a metal, sulfur, a sulfide, carbon, or non-volatile organic
matter, such as tartaric acid, citric acid, sugar, starch, etc., must be
present. If non-volatile organic matter is present, carbonization
will take place on warming.

SOz with separation of sulfur indicates the presence of a sulfo-
cyanate, in case there was no action with dilute sulfuric acid.

CO from oxalates and other organic substances, and cyanates.
It is an odorless gas, which does not fume in the air and burns with
a blue flame. |

(b) Colored.

Cl, a yellow gas with a suffocating odor, turns iodo-starch paper
blue, and indicates the presence of both a chloride and an oxidizing
substance.

430 SYSTEMATIC ANALYSIS

ClO2, a yellow gas, very similar to chlorine, but which explodes
violently on being heated, indicates a chlorate. If the substance —
deflagrates on being heated on charcoal, only a small portion of the
substance should be used for the test with concentrated sulfuric acid;
but if no explosion takes place on warming, more of the substance
should be added.

HBr from bromides has a penetrating odor, fumes in the air, and is
always colored yellowish brown by the presence of small amounts
of bromine. The sulfuric acid is at first colored brown in the case
of a colorless bromide, but becomes colorless on being boiled.

CrO2Clz, brown (similar to bromine), results from the presence of
a chloride and chromic acid.

Iz, violet. In the case of a colorless iodide, the sulfuric acid is
at first colored brown by small amounts of iodide, or gray, solid iodine
is deposited if considerable iodide is present, which volatilizes on
warming, forming violet vapors. If considerable iodide is used for
this test, the sulfuric acid is reduced to SOz, or even H2S (ef. p. 301).

Mn207, violet, is formed from permanganic acid, and is devom-
posed with scintillation, often exploding, on being warmed.

NOz, brown, with a.penetrating odor, comes from nitrates.

After the preceding tests have been made, the next step is the

Solution of the Substance

As solvents the following are used:

1. Water;

2. Hydrochloric acid;

3. Nitric acid;

4. Aqua regia.

In the majority of cases the first three solvents suffice, aqua regia
being seldom necessary, as will be seen from the following table:

SUBSTANCES SOLUBLE IN WATER

Of Group I (p. 284) the following are soluble:

1. Chlorides.—All except AgCl, CusCle, HgeCle, PtCle, AuCl,
BiOCl1, SbOCl1, Mg2OCle. PbCle and TIC! are difficultly soluble.

2. Bromides.—The same as the chlorides.

3. Iodides.—All except AgI, HgeI2, Hgle, Cusl2, PdIz, Til; PbI2
is very difficultly soluble.

4, Cyanides.—Only the cyanides of the alkalies, alkaline earths,
and mercury.

SOLUTION OF THE SUBSTANCE 431

5. Ferrocyanides.—Only those of the alkalies and alkaline earths.

6. Ferricyanides.—Same as the ferrocyanides.

7. Cobalticyanides.—Only those of the alkalies, alkaline earths,
and. the ferric, mercuric, and lead salts.

8. Thiocyanates.—Those of the alkalies, alkaline earths, iron,
cupric copper, and mercuric mercury.

9. Hypochlorates.—All.

Of Group II (p. 284) the following are soluble:

10. Nitrites—All. Silver nitrite is difficultly soluble.

11. Acetates.—Silver and mercurous acetates and certain basic
acetates are difficultly soluble.

12. Cyanates.—Those of the alkalies, alkaline earths, and most
of the remaining ones. Silver and lead cyanates are insoluble.

13. Sulfides—Only those of the alkalies and alkalines earths.
CaS is difficultly soluble.
14. Hypophosphites.—All.

Of Group III (p. 284) the following are soluble:

15. Sulfites —Those of the alkalies, and the bisulfites of the alkaline
earths.

16. Carbonates.—Those of the alkalies, and the bicarbonates of
Ca, Sr, Ba, Mg, Fe, Mn.

17. Oxalates.—Those of the alkalies; the remainder are diffi-
cultly soluble or insoluble. Most oxalates, however, with the exception
of Ba, Ca, and Sr oxalates, form soluble complex salts with alkali
oxalates. |

18. Iodates.—Only those of the alkalies.

19. Borates.—Those of the alkalies. The remaining borates are all
difficultly soluble in water, but soluble in ammonium chloride as a rule.

20. Molybdates.—Only those of the alkalies.

21. Selenites.—Those of the alkalies are readily soluble, the re-
maining ones are difficultly soluble.

22. Selenates.—All except the barium and lead salts.

23. Tellurites.—Only those of the alkalies.

24. Tellurates.—Only those of the alkalies.

25. Tartrates.—The normal tartrates of the alkalies, and lithium
and sodium bitartrates. The remaining tartrates are insoluble in
water, but are usually soluble in an excess of alkali tartrate solution,
forming complex salts.

26. Citrates.—Only those of the alkalies are readily soluble in
water. The insoluble citrates usually dissolve in an excess of alkali
citrate solution.

but more soluble in acids;

stances, small letters for the unimportant ones.

i,-slightly soluble in water and slightly soluble in acids;
The salts referred to are generally consideféd™

432 SYSTEMATIC ANALYSIS
SOLUBILITY
ie: 2, 34 se | 8, | ac let at] + i af
Os } Wiw;wiw|] Wi wiws|wiwl|wiwil wl wa
ak aN } Wiwiwi|alatftalatatsltatatataia
aS \ Wiwiwia a a|jajsj—|—|—latstatla
Borate, BC;=..| W | W | w |w-a| a aj/al|as}a|a|s |) eae
* Bromide, Br~..|W|Wi]wjf})wt{|wiifiwtswti]w {\|wki| w| wl] wlw
Carbonate, Viwiwiw|a|alajal—|—|al|a]|a]a
Chlorate, ClO;-|} W| w|wi|w!]WwWfiwtiw wi|wi|wl|wl|w
Chloride, ClI-...; W| W|W|W!] W|W|W well W| W| WwW) WwW |
eerie } Wi WwWiwi wi w-alwe-a} a ==). | W |
Cyanide, CN.... W | W| w] wiew | wiw-al —|a|Ada J ai} ai
MeCN ee } WiwWiwiwiw * wil—|—l|a I I I
¥ [Fe ie Wiwiwiw|w w-a| —|—|A-I| a jor
Fluoride, F~ w | w | W | a-1| A-I | w-al| w-a| a |w-a| a | w-a| w-a| w-a
Hydroxide, \)wilwilwi|alwalwi/wi/a|a/al|alala
Iodide, I~..... wiwiwiw| WwW |wi|wiw w|Wwiwi]w
Nitrate, NO;~ | W Ww Wiw|wiwiwiwiwiwiwiwiw
Oxalate,C.0.-|W/}W;}W|]aj|A | alaslai|w-al a |w-al ala
Oxide, O..... wilwf{]—]A{W-A| W/] WiA-T;A-T}] A} A} AIL A
eae lwiwi/wi/alalalaAlaAlal/alAalAala
Silicate, SiOZ--| W | W| —]| a a ajajA-I}a/aljafatsa
Sulfate, SO.-...;W | W|W|WyW-l| I] I | W/W-ll WwW} ws) wi w
Sulfide, S=..... Wi|wiwi!]a{f{w-A|Wiwisajfjai]/A;}|A!]AIA
Peewee, } Wiwiwiwis]wfwl|w wiwiwi{|wiw
Shree ke } W|W!]W {veal a a | a w | a |w-al a’| W
W or w, soluble in water; A ora, insoluble in water, soluble in HCl, HNOs or aqua regia;

A-I

*Based on the table in Fresenius-

SOLUTION OF THE SUBSTANCE 433
‘TABLE *
ies seek, ea fc: Oe ea Bee

ed aldlGlelegibealee

wi|Wi w | W{|w-al w |W] w |] wi wl] w | — |—|— eee
Sie e@ 18 | a f ala} a | a |—|;—)—{—|— ee
ala} ajaj|jla|ai4|ala {a |j|—|—|—s|—|— wer.
ajal| a aij—|— /]a} a _|w-a} a |] — | — |—|— |Borate, BO=
wiw) i |w-il ai} w | w/! w-a| W |.w|w-a|w-a| w | w |Bromide, Br~
Al—| a |Alal]aJ{A}] aja |—|—|—|—]/— { ae
wiw!| wiwiw|wiiw| w]wf|wi|— |] — |—|— (Chlorate, ClO;—
W I |W-I| A-I[} W |W|W-A| W W |W-A| W | W |Chloride, Cl—
—|w| a |A-I| a | w-ajw!] a SLE ep ae = Ta ge es
a-ij|—j; I a!—|} Wia}j|— ]aj—j|—l— w |Cyanide, CN
Boe i ee) | fj) i pedi) | — 3 (TFeCN)
E}t] i jaj—j;—j1| i faijil i |—|-|—|{ Pea
w-alw| a | @| wj|wajla| w |w-alw| w | w |—|—|Fluoride, F~
Aja|—|a|—|—|a| a] a l—|a| a |—|—|{Bycronde,
Wiw; I |wW-lI| A} A |W| a | W/W w {|w-a| a | i |Iodide, I-
Wiw; W/|Wi|w Ww |W! W-A| W }—| — | — |—| w Nitrate, NO;-
og 3) EE oe a a ala Sia when tt Oxalate, C,0.=
Aj}A} A j|A|{A{A {A} A | A{AJA-I] A |—] a |Oxide, OF

A Al 2 |} Aj}aya {}A} a+a|a| A fw-al—|— Rape ey
aja}—}|ad—}|—]/a|j— }]a |—| — | — |—|— Sikcate, Si0,-
W/W W-A}|A-I| W | W-A|W| w-a | W |W} — | — |—| w Sulfate, SO.=
Bee See ACE ASIA) OA (AIA A |A| A |Sulfide, S—
wiw| I aj};Al|W |-a}| — |w-al— —=(—1 & Bors ees
w-a|}W| a & |w-al a@ | w| a |w-ala|—|a |—|— Ree

I or i, insoluble in water and in HCl, HNOs or aqua regia;
or a-1, insoluble in water, slightly soluble in acids.

Capitals

as normal. Acid and basic salts are omitted in this table.
Wells Qualitative Analysis.

-A or w-a, slightly soluble in water

are teed TOF th

or the more important sub-

434 SYSTEMATIC ANALYSIS 4

27. Pyrophosphates.—Only those of the alkalies.
28. Metaphosphates.—Only those of the alkalies.

Of Group IV (p. 284) the following are soluble:

29. Phosphates.—Only those of the alkalies.

30. Arsenites.—Only those of, the alkalies.

31. Arseniates.—Only those of the alkalies.

32. Thiosulfates—Almost all are soluble, though the silver and
barium salts are difficultly soluble.

33. Chromates.—Those of the alkalies, Ca, Sr, Mg, Zn, Mn, Fe,
and Cu are soluble, the others are difficultly soluble or insoluble.

34. Vanadates.—The orthovanadates are unstable; the pyro-, meta-,
and polyvanadates are soluble in water, as a rule. The lead and
mercurous salts are insoluble, also the vanadates of the iron group.

35. Periodates.—All more or less soluble in water, except sliver
periodate, which is insoluble.

Of Group V (p. 285), the following are soluble:

36. Nitrates.—All except a few basic salts.

37. Chlorates.—All.

38. Perchlorates.—All.

39. Manganates and Permanganates.—All.

Of Group VI (p. 285), the following are soluble:

40. Sulfates.—All except the Ca, Ba, Sr, and Pb salts, and a few
basic sulfates.

41. Fluorides.—Those of the alkalies, silver, ae mercury; the
remaining fluorides are difficultly soluble or insoluble in water.

Of Group VII (p. 285), the following are soluble:

42. Silicates.—Only those of the alkalies.

43. Tungstates.—Only those of the alkalies.

Of the salts insoluble in water, all dissolve in acid (hydrochloric or
nitric) except AgCl, AgBr, AgI, AgCN, AuCl, PtCle, BaSO4, SrSOa,
PbSO4, HgS, Prussian blue, CaF2, SnSe (mosaic gold), SiOz, many
silicates, fused PbCrOxz, and the strongly ignited oxides: AleOs, CreOz,
TiO2, SnO2z, SbeO3.* TiOes, SnOe, and Sb2O3 can be dissolved by
long continued boiling with concentrated hydrochloric acid.

Of the salts insoluble in acids, the following dissolve in aqua regia:
PtCle, AuCl, HgS, SbeOzs, SnSe, and Prussian blue (after long treat-
ment).

The following substances are not dissolved by aqua regia: AgCl,

* The oxides of antimony are changed to Sb.O, after long ignition in the
alr.

SOLUTION OF THE SUBSTANCE 435

AgBr, AgI, AgCN, BaSO., SrSO4, PbSOs, CaF2,* fused PbCrO,,
AlzO3z, CreO3, native TiO2 (rutile, anatase, brookite), native SnO2g
(cassiterite, tinstone), SiOz, Si, many silicates, C, carborundum, and
strongly ignited iridium (rhodium, ruthenium, and osmium).

In order to bring such substances in solution it is necessary to
subject them to a special treatment. The process to be chosen depends
largely upon the nature of the insoluble substance, so that a few general
tests are necessary before going farther. Very often the preliminary
examination will have been sufficient, but it is always well to perform
the following simple experiments:

1. Heat a small portion of the residue insoluble in acids on the
charcoal stick to see whether a metallic button can be produced.

(a) No metallic button is produced. ‘The absence of silver, lead,
and tin is thereby assured.

(b) A metallic button is formed. The button is flattened in an
agate mortar, and its solubility in acids is tested.

(a) The metal dissolves in nitric acid forming a clear solution,
showing the absence of tin. Add a little hydrochloric acid to the nitric
acid solution; a curdy precipitate is formed if the metal is silver, con-
sisting of silver chloride, insoluble in water, but soluble in ammonia.

If the nitric acid solution becomes turbid on the addition of sul-
furic acid, lead is present.

(8) The metal does not dissolve in nitric acid forming a clear
solution, but leaves a white, insoluble powder: metastannic acid.
Treat a new button with concentrated hydrochloric acid, when it will
completely dissolve if silver is absent. Mercuric chloride produces
a white precipitate of mercurous chloride in the hydrochloric acid
solution: tin is present.

2. Heat a second portion of the insoluble residue in a small test-
tube with concentrated sulfuric acid and test to see whether the esca-
ping gas renders a drop of water turbid.

A turbidity shows the presence of an insoluble fluoride (CaF 2).

3. Heat another portion of the residue (with the help of a platinum
_ wire) in the upper reducing flame of the gas-burner, allow to cool in
the inner mantle, moisten with dilute hydrochloric acid, and notice
whether the odor of hydrogen sulfide can be detected. Then test
to see whether it will now impart a characteristic coloration to the
flame. The presence of a sulfate is betrayed by the odor of hydro-
gen sulfide, and the flame test shows whether barium alone or a mix-
ture of barium, calcium, and strontium is present.

* Calcium fluoride will be dissolved by the long continued action of aqua regia.

436 SYSTEMATIC ANALYSIS

4. Test another portion of the residue in the salt of phosphorus
bead; silicic acid or a silicate usually gives a skeleton bead (cf: p.
419).

As the skeleton bead is not always obtained even when silica is
present, a further test for silicic acid is often necessary (cf. p. 418).

5. Now heat the salt of phosphorus bead in the reducing flame
to test for the presence of titanium, which causes the bead to become
violet. The violet color appears more quickly on the addition of a
little piece of tin-foil. If iron is present at the same time, as is always
true in the case of rutile, the bead is colored brownish red in the redu-
cing flame.

6. The presence of chromium is often detected by the green color
of the residue. In the case of chromite (gray or black residue) fuse
some of the substance with sodium carbonate and potassium nitrate in
the loop of a platinum wire (cf. p. 142), when a yellow melt is
obtained if chromium is present, which (after being dissolved in
water and acidified with acetic acid) yields, with silver ee
reddish-brown precipitate of silver chromate.

7. If the residue is gray or black, it may also consist of carbon.
Heat a small portion upon a piece of platinum-foil; if carbon is present,
the mass will glow and, if it does not burn completely, a lighter-
colored ash will be obtained. In doubtful cases melt a little potas-
sium chlorate in a test-tube, and add a little of the insoluble residue,
when a distinct glowing or a little explosion will take place if carbon
is present. It is necessary to avoid the addition of shreds of filter-
paper in this test.. _

8. Silicon and Silicides (carborundum, etc.) are seldom met with,
and show the greatest stability toward the above-mentioned reagents.
By fusing with caustic alkali in a silver crucible, however, they are
readily decomposed with evolution of hydrogen (cf. p. 421).

After dissolving the melt in water and acidifying, gelatinous silicic
acid separates out, particularly after evaporation.

MetuHops FOR GETTING SUBSTANCES INTO SOLUTION WHICH ARE ~

INSOLUBLE IN ALL ACIDS

1. Insoluble Halogen Compounds (the silver compounds alone
come into consideration) can be brought into solution by melting
the mass, adding a little dilute sulfuric acid and a piece of zine
so that it comes in contact with both the acid and the insoluble
substance. After a while pour off the acid; it contains the halogen
acid in the presence of zine sulfate, and should be kept for the sub-

Te eee eh Te ee a ee ee a ee

a aan
TS SS he ee

SOLUTION OF THE SUBSTANCE 437

sequent tests for acids, cf. p. 306. The residue consists of metallic
silver. Wash it with water, dissolve in dilute nitric acid, filter and
test the solution for silver with hydrochloric acid.

2. Insoluble Sulfates of the Alkaline Earths are brought into solu-
tion by fusing in a platinum crucible with four to five times as much
calcined sodium carbonate, or with a mixture of equal parts of sodium
and potassium carbonates. Mix the finely powdered substance
in the crucible with the sodium carbonate, cover the mixture with a
thin layer of more carbonate, place the lid on the crucible and heat
at first gently over a small flame in order to drive off the moisture
which the carbonate always contains, and then raise the temperature
until the mass fuses to a thin liquid; maintain this temperature for
about fifteen minutes. Remove the fused mass from the crucible
as directed on p. 418. Heat with a little water on the water bath
until the fused mass disintegrates, and no more hard lumps can be
felt with a glass rod, then filter. The filtrate will contain the sulfate
as sodium sulfate, and the residue will consist of carbonates of the
alkaline earths. Wash it with a 5 per cent sodium carbonate solution
until no more sulfuric acid can be detected in the filtrate, and then
wash with hot water until the wash-water no longer reacts alkaline
(cf. p. 109). Dissolve the residue in nitric acid, and analyze as
described on p. 110.

3. Lead Sulfate may be boiled with a concentrated sodium car-
bonate solution, which forms insoluble basic lead carbonate and
soluble sodium sulfate; with caustic soda, which forms soluble lead
plumbite and sodium sulfate; or with ammonium acetate (cf. p. 209).
Calcium sulfate is also decomposed completely by boiling with soda
solution, as is strontium sulfate (though less readily), but barium
sulfate is incompletely decomposed.

4. Silicic Acid and Silicates should be fused with sodium ae
ate, exactly as described on p. 418.

5. Metastannic Acid, as obtained by the oxidation of tin with
nitric acid, is readily dissolved by boiling with a little concentrated
hydrochloric acid, and then treating with considerable cold water
(ef. p. 257).

Tin dioxide, as it occurs in nature (tinstone), as well as the strongly
ignited metastannic acid, cannot be brought into solution in this
way. One of the methods mentioned on p. 260 (usually the
sodium carbonate and sulfur method) must be used.

6. Insoluble Fluorides are first heated with concentrated sulfuric
acid, and the sulfate formed is brought into solution by the method
described under 2.

438 SYSTEMATIC ANALYSIS

7. Titanium: Dioxide is fused with potassium pyrosulfate in a
platinum crucible (cf. pp. 180 and 158); or it is fused with sodium
carbonate, the melt treated with cold water, and the residue dissolved
in hydrochloric acid (ef. p. 157). Heating for some time with con-
centrated sulfuric acid will dissolve pure titanium dioxide. When
cold the solution may be diluted.

Fusion with potassium pyrosulfate is also suitable for decomposing
native aluminium oxide (corundum).

8. Chromium Trioxide and Chromite are fused with sodium ear-
bonate and a little potassium nitrate in a platinum crucible or with
sodium peroxide in a nickel or iron crucible (cf. p. 133).

9. The Insoluble Complex Cyanides are completely decom
by boiling with caustic soda in a porcelain dish.

After boiling with the alkali, dilute with water and filter. The filtrate
will contain the acid in the form of its sodium salt; and, in some cases, may
also contain aluminium and zinc. Saturate the filtrate with carbon dioxide,
boil and filter off any precipitate (Al(OH); or ZnCO;); dissolve this precipitate
in hydrochloric acid and test for zinc and aluminium. Acidify the alkaline
filtrate obtained above with hydrochloric acid and test for ferrocyanic and
ferricyanic acids according to pp. 318 and 321.

The soluble complex cyanides are decomposed before the analysis by

heating them with concentrated sulfuric acid (cf. p. 153),

REACTIONS THAT ACCOMPANY THE DIsSSOLVING PROCESS

When a substance is dissolved, whether in water or in acids,

phenomena are often observed which may be of great importance
2s concerns the subsequent analysis. Moreover, the color, reaction
of the solution towards indicators, or the evolution of gases will lead to
important conclusions. First, test the substance with regard to its solu-
bility in water, by taking about 0.5 gm. of the fine powder, adding a
little cold water, and noting whether any bubbles of gas are given off.

A gas is evolved when there are present:

(a) Peroxides of the Alkalies or Alkaline Earths, which are partly
decomposed into hydroxide and oxygen:

2Na202+2H20 =4Na0H+02 t ;
2Ba02-+2H20 = 2Ba(OH)2+0z2 1.

Barium peroxide is decomposed in this way only by heating the water.

Test the escaping gas for oxygen by means of a glowing splinter.

In the alkaline solution (red litmus is changed to blue) some unde-
composed peroxide will still be found.

Dilute the solution with considerable water, cool, and carefully acidify with
sulfuric acid, add a little ether, some potassium dichromate solution, and

: ’
(a a,

SOLUTION OF THE SUBSTANCE 439

shake the mixture. If a peroxide is present, the upper ether layer will now
be colored blue. A better method for detecting the hydrogen peroxide, formed
by the action of the sulfuric acid upon the peroxide, consists in adding a few
drops of titanium sulfate solution; a distinct yellow color will be noticed if
- only traces of hydrogen peroxide are present (cf. pp. 84 and 159).

(b) Carbides of the Alkaline Earths (calcium carbide).
These are decomposed into acetylene (which has a peculiar odor,
and burns with a luminous flame) and calcium hydroxide:

CaC2+2H20 =Ca(OH)2+Ce2He t -
(c) Nitrides of the Alkaline Earths (magnesium nitride).

Magnesium nitride is decomposed by water into magnesium
hydroxide and ammonia:

Mg3N2+6HOH =3Mg(OH)2+2NHs3 ff.

| If considerable water is added, there is no gas evolution, because

the ammonia will be absorbed by the water; ‘but on boiling the solu-
tion, ammonia will be given off, which can be readily recognized by
its odor.

' (d) Phosphides of the Alkalies and Alkaline Earths.—These are
decomposed by water, setting free spontaneously combustible phos-
phine:

CazP2+4H20 = Pe2H4 tT +2Ca(OH)>.

Very small quantities of the phosphide can be recognized by the
characteristic garlic odor.

(e) Many Chlorides, Bromides, and Iodides of the Negative Ele-
ments, e.g., PCl3, PCls, etc., are decomposed into the halogen hydride
and the oxygen acid of the negative element:

PCl5+4H20 =5HCI+H3POs,.

(f) A few Sulfides which are Seldom Met with (Mg§, AloSs, etc.).—
These are decomposed by water with loss of hydrogen sulfide, which
can be detected by its odor, and by its blackening lead acetate paper:

MgS+2H.0=Mg(OH)2+H28 } .

After any reaction caused by the first addition of water is over,
add about 10 to 15 ec. more, heat the water to boiling and then
allow it to cool.

If the substance dissolves completely, forming a clear solution,
it is evident that it is unnecessary to test for any insoluble substances
in the subsequent analysis.

440 SYSTEMATIC ANALYSIS

If a residue remains, it is possible that a part of the substance —
has dissolved in the water. To determine whether this is the case,
decant the liquid through a filter and carefully evaporate a little of
the filtrate to dryness on platinum-foil (or a watch-glass). If the foil —
is heated too hot, volatile compounds may escape unnoticed. If a a
residue remains after evaporation, it is evident that a part of the —
original substance is soluble in water. Then treat the original residue —
several times with small amounts of water, and analyze the aqueous —
extract thus obtained by itself. Treat the part remaining undissolved
with acid, using hydrochloric acid unless the preliminary examina-
tion has shown the presence of either lead or silver, when nitric
acid should be used.

Treat the residue with 0.5—1 cc. of 12-normal acid (notice whethael
there is any evolution of.a gas), heat gently, and then dilute with — |
water, to dissolve any chlorides insoluble in hydrochloric acid. It —
must be remembered, however, that bismuth and antimony salts —
form insoluble basic chlorides on dilution with water, so that too —
much water should not be added. :

If a residue remains after treatment with acid, bring it into solu- :
tion by one of the methods described on pp. 436-8.

II. EXAMINATION FOR THE METALs (CaTIONS) a

TaBLE XI.—GENERAL SCHEME FOR SEPARATING THE METALS INTO GROUPS ~ _

~*~

Solution may contain all the common basic constituents. Add HCl in slight” %
excess. (1) .

Precipitate: | Filtrate: Groups II, II, IV and V. Saturate with HS. (8).
Group I. F
Examine asout-| Precipitate: | Filtrate: Groups III, IV and V. Test for “4
lined in Table Group II. phosphoric acid. If found present modify the q
X, p. 282. (2) | Examine as| following procedure as indicated in Table XII.
outlined in| Add NH,OH and (NHs4) Ss. (5) ;

Table VII, p.
272. (4) Precipitate: | Filtrate: Groups IV and V:
Group III. Add (NH 4)2CO3. (7) J
If phosphate is “-
absent exam-| Precipitate: Filtrate:
ine as outlined | Group IV. Group V.

“7

in Table V,p.| Examine as| Examine as
189 orinTable | outlined in| outlined in
VI, p.192. If | Table III, p.| Table I, p.
phosphate is| 111 or in Ta-| 97 or in Ta-
present exam-| ble IV,p.113.| ble II, p. 99.
ine by Table} (8) (9) Fal
XII, p. 445. .
(6) . .

EXAMINATION FOR THE METALS 441

- PROCEDURE

1. Add HCl as directed on p. 283. If the original substance was com-
pletely soluble in dilute HCl, it is evident that no silver or mercurous salt is
present. Often, when lead is present, the solution is clear while hot, but lead
chloride is deposited as the solution cools. It is usually best to filter off
such a precipitate, but it will be changed to less soluble lead sulfide when H.S
is introduced to precipitate the second group. If the original solution is
alkaline to phenolphthalein- or to litmus, a precipitate may form when none
of the metals of the first group is present. Thus a solution of sodium silicate
gives a white, gelatinous precipitate of silicic acid, a solution of an alkali tung-
state gives a precipitate of tungstic acid, a solution of a thio salt of arsenic,
antimony or tin gives a colored sulfide precipitate and the solution of a complex
cyanide may form a precipitate when neutralized. These precipitates, however,
are not likely to be mistaken for a chloride of silver, lead or mercury. If-a silicate
is present, it is absolutely necessary, to remove the silicic acid at the start
by the method given on p. 416 and tungstic acid may be removed in exactly
the same way. If a thio salt is present, examine the precipitate according to
Table IX, p. 275, and test for alkalies and alkaline earths according to Table
Il], p. 111 and Table I, p. 97. If the original solution is alkaline, it is
necessary to test for iron and aluminium only when the solution contains non-
volatile organic matter which prevents the precipitation of these elements by
hydroxyl ions. The addition of HCl may cause the precipitation of BiOCl
and dilution may cause the precipitation of BiOCl, SbOCI or a basic salt of
some other metal, especially titanium and tin. With the exception of the
titanium precipitate, such basic salts are easily dissolved by filtering and
treating with 6-normal HCl, or the basic salts of antimony, bismuth and tin
may be changed into less soluble sulfides by introducing H.S without filtering.

2. Examine the precipitate of Hg,Cl., PbCl, and AgCl exactly as described
on p. 283. For the detection of thallium, which is sometimes precipitated
with this group, consult Part V.

3. Transfer the filtrate from (1) to a 300 ec. Erlenmeyer flask, fit the latter
with a two-holed rubber stopper and insert through one hole a right-angled
glass tube which reaches nearly to the bottom of the flask and through the
other hole a shorter tube, similarly bent, which reaches only to the bottom
of the rubber stopper. Raise the longer tube till it is just above the surface
of the solution, heat the solution to boiling and begin passing a steady stream
of H.S through the longer tube. -Remove the flame from beneath the flask,
close the shorter tube with a piece of rubber tubing which has one end stopped
up with a short piece of stirring rod, and lower the longer tube so that it dips
well below the surface of the solution. Shake well and continue keeping up
the pressure of the H:S; in this way the gas is absorbed as fast as the sulfide
precipitates and the solution is kept saturated with the gas without wasting
a great deal of it. When the precipitation appears to be complete, shut off the
HS, open the flask, add an equal volume of cold water and again saturate with
H.8. Close the short tube and shake the flask well for two or three minutes
while keeping up the H.S pressure. Finally filter off the precipitate and wash
it. promptly with H.S water. If an oxidizing agent is present,* considerable

* If much oxidizing agent or considerable arsenic acid is present, it is best to
pass SO, into the hot solution until a complete reduction is accomplished and then

442 SYSTEMATIC ANALYSIS

free sulfur will be deposited and this will greatly delay the precipitation of the
sulfides. In case it is desired to know whether a precipitate contains nothing
but sulfur, wash it several times with alcohol, then with carbon disulfide (away
from any free flame) and then with alcohol again; this treatment will serve to
remove the sulfur. It is important to adjust the concentration of the acid
properly before introducing H.S by measuring the quantity added in getting
the substance into solution and in precipitating the first group. After diluting
with water, as above directed, the solution should be about 0.3-normal in acid,
if more concentrated cadmium and lead will not precipitate and if less con-
centrated sulfides of zinc, nickel and cobalt may precipitate. On the whole
it is better to err with too little acid rather than with too much, as enough zine,
nickel and cobalt will always remain in the filtrate to give a test in the next

group and the presence of these elements does not seriously interfere with the ~

analysis of the second group. Like cadmium, zine gives a white ferrocyanide
in the confirmatory test for copper and nickel gives a faint blue with ammonia;
either nickel or cobalt will interfere with the final test for cadmium, but the
treatment outlined on p. 275 will overcome this difficulty. If after pre-
cipitating with hydrogen sulfide a turbid filtrate is obtained, due to free sulfur,
prepare some filter-paper pulp by shaking pieces of filter paper in a bottle
with hot water, add some of the pulp to the filtrate and filter through a fresh
filter. The hydrogen sulfide precipitate oxidizes somewhat on being exposed

to the air, and a little soluble sulfate is likely to form which forms a precipitate —

on coming in contact with H.S in the filtrate. For this reason the precipitate
should be washed promptly with hydrogen sulfide water without ever letting
the filter drain completely until the washing is finished. If the filter clogs,
place the filter and precipitate in a beaker, shake it up with hydrogen sulfide
water and filter through a fresh filter. In qualitative analysis, all but the first
washings of a precipitate should be discarded, as a rule.

4. Examine the hydrogen sulfide precipitate as directed on p. 273. If
gold, platinum or considerable tin is present, these elements are often found
in the residue of mercuric sulfide obtained after treatment with nitric acid.
When the presence of these elements is suspected, take a little of the residue
for the mercury test and fuse the remainder in a porcelain crucible with a
mixture of equal parts potassium cyanide and sodium carbonate. Cool, wash
out all the soluble alkali salts with water and discard this solution. Gold,
platinum, tin and lead will be left behind in the metallic condition. Treat
the metallic residue with dilute nitric acid and test the solution for lead in the
regular way with sulfuric acid (p. 274). Heat the residue of gold, platinum

and metastannic acid with concentrated hydrochloric acid. Dilute, filter and

test for tin with HgCl, in the usual way (p. 276). Dissoive any gold or platinum
in aqua regia, add ammonium chloride, evaporate to dryness on the water
bath and treat the residue with a very little water; a yellow precipitate of
(NH,)2PtCl, shows the presence of platinum. Filter, test with FeSO, for
gold, and confirm by the charcoal stick reaction (p. 266).

5. Take a little of the filtrate from (3), boil off the hydrogen sulfide, add
a little bromine water to oxidize any iron and the last traces of hydrogen
sulfide and make alkaline with ammonia, If a precipitate forms it may con-

remove the excess of SO, by a stream of CO.. If the excess of SO, is not removed
it reacts with H.S as follows: 2H2S+SO.>5 2H,0+38.

eee

-“j
;
“4
a
Pe.
‘
a
.

ANALYSIS IN THE PRESENCE OF PHOSPHORIC ACID 443

sist of a phosphate of barium, strontium, calcium or magnesium, or an alkaline
earth fluoride or oxalate. Phosphoric acid very often occurs in minerals, and
for this reason a special procedure is often required for the analysis of Groups
III and IV. Without stopping to filter off the precipitate produced by
ammonia, dissolve it by the careful addition of a little nitric acid, heat the
solution nearly to boiling and add an equal volume of ammonium molybdate
reagent. A yellow precipitate, which may be slow in forming, shows that
phosphoric acid is present. Arsenic acid gives a similar precipitate (p. 231),
but arsenic should not be present at this stage in the analysis. If phosphoric
acid is found present, examine the precipitate as outlined in Table XII. If
an oxalate or fluoride is indicated by the preliminary examination, especially
by the test with concentrated sulfuric acid, it is best to remove these acids
by heating the original substance with concentrated sulfuric acid. This is
likely to leave, after diluting, an insoluble residue of alkaline earth sulfate
which should be examined as directed on p. 437. If the behavior of the
original substance in the closed tube test indicated the presence of non-volatile
organic matter, it is necessary to remove it before proceeding with the analysis
of Group III, because tartaric and citric acids, sugars, starches and similar
substances prevent the precipitation of iron, aluminium and chromium with
ammonia. Such organic substances can be removed by ignition or by. repeated
treatment with concentrated sulfuric and nitric acids: adding about 5 cc. of
sulfuric acid, an equal volume of-concentrated nitric acid and evaporating till
strong fumes of sulfuric acid are evolved, cooling and repeating the treatment
with nitric acid as often as necessary. The ignition treatment often makes
Al,O;, CrsO3, Fe2O; and SiO, insoluble. Fuse such a residue with KHSO, and
examine it by itself. The treatment with sulfuric and nitric acid is likely
to leave an insoluble sulfate behind; fuse it with sodium carbonate (p. 109).

6. If phosphate is found present, analyze as outlined in Table XII, p.
445. If phosphate is absent, analyze Group III by method A, p. 189 or by
method B, p. 192. Many of the elements in this group commonly occur in
different states of oxidation. In reporting the final results of the analysis it
is not sufficient to state that iron, chromium, or manganese is present, but it
should be stated in what condition such an element is present in the sample
as received. It is necessary to determine this by special tests, using the
characteristic reactions described in Part II under the element in question.

7. Treat the filtrate from (5) with (NH,),CO; according to Method A, p.
111 or according to Method B, p. 113. In the former case magnesium will
not be precipitated with this group and in the latter case it will be precipi-
tated. In most cases, Method B will be found preferable.

8. Examine the (NH,),CO; precipitate as directed on p. 111 or as directed
on p. 113.

9. Examine the alkali group according to the directions on p. 97 or on
_ p. 100, omitting the magnesium test provided Method B was used for the
analysis of Group IV. ‘Test a portion of the original substance for ammonium.

ANALYSIS IN THE PRESENCE OF PHosPHORIC ACID

There are three classic methods for examining a solution for the members
of Groups III and IV when phosphoric acid is present. The first of these
methods depends upon the fact that when tin is boiled with strong nitric acid,

4d : SYSTEMATIC ANALYSIS

insoluble metastannic acid is formed (p. 257) and this precipitate carries down
with it phosphoric acid, arsenic acid and to a lesser extent other substances —
such as ferric oxide, titanium oxide, etc. It is probable that an adsorption —

compound is formed, rather than a simple chemical compound. Instead of ‘-

using metallic tin, it has been found possible to accomplish the same result “4
by preparing metastannic acid in advance and adding it to the solution.*

After the phosphoric acid has been precipitated in this way, it is filtered off s

and the filtrate examined for Groups III and IV in the usual manner. “7
The second method of analysis in the presence of phosphoric acid is the
so-called basic acetate method. If a solution containing ferric iron or aluminium __
is carefully neutralized, ferric phosphate or aluminium phosphate is precipitated
before an insoluble hydroxide or an alkaline earth phosphate is formed. If
the neutralization goes too far, and an excess of hydroxyl ions is provided,
ferric and aluminium hydroxides are likely to be formed and alkaline earth
phosphate will then precipitate. One of the best methods of preventing the
neutralization going too far is to boil the dilute solution with a little acetic
acid and considerable sodium acetate. Then, if a slight excess of iron is
present, all of the phosphoric acid will be precipitated as ferric phosphate and
the excess of iron will be precipitated as basic acetate. It is better to use fer- —
ric iron rather than aluminium in this separation, because basic ferric acetate
by its color shows when enough trivalent metal is present. a
The third method of analysis is the bariwm carbonate method. This method
is the same in principle as the basic acetate method except that the solution —
is neutralized with barium carbonate in the cold. The objection to this method —
of analysis is that an alkaline earth is added so that a separate portion of the
solution must be used in the tests for this group. In some special cases calcium

carbonate, zinc oxide and cadmium carbonate are used instead of barium

carbonate. All of these substances will neutralize a solution sufficiently to —

precipitate phosphates and hydroxides of iron, aluminium and chromium. They —

do not, however, precipitate bivalent metals as a soluble carbonate would..

The barium carbonate method is useful for separating the trivalent from — ;

the bivalent metals of Group III and will be described in Volume I of this
book. The other two methods will be outlined here.

PROCEDURE. TIN METHOD

1. Evaporate the filtrate from Group II just to dryness. Add 10 ce. of +
concentrated HNO;, evaporate to dryness and repeat the evaporation with

HNO; once more. Finally add 10 cc. of concentrated HNO; and introduce —

about 1 gm. of tin foil in small portions. Boil to small volume in order to com-
plete the precipitation of the tin and pour the concentrated solution into —
100 cc. of water contained in a narrow cylinder such as a 100-cc. graduate. —

Next morning siphon off the supernatant liquid and discard the precipitate

of metastannic acid, which should contain all the phosphoric acid. s
2. The tin foil usually ‘contains traces of lead and copper. To remove —
these, saturate the solution with H.S and filter. Reject the precipitate. §
3. Treat the precipitate with NH,OH and (NH,).S, examine the precipitate _
for Group III and the filtrate for Groups IV and V exactly as outlined in —
Table XI. |

*Cf. W. MeckLenBura, Z. anal. Chem. 52, 293 (1913)

BASIC ACETATE METHOD 445

TABLE XII.—ANAtysis oF Groups III anp IV 1n PRESENCE or PHOSPHATE

Tin Method. Basic Acetate Method.

Remove HS and HCl from filtrate from | Treat filtrate from Group II with NH,OH
Group II, and evaporate repeatedly with | and (NH4)S. Filter and keep filtrate.
HNO;. Add tin foil, concentrate, dilute | Examine the precipitate according to Table
and allow the precipitate to settle. (1) VI up to and including treatment with

. HNO; and KCI0O;3. Test for Fe+++ in
filtrate from MnO, precipitate. Add
NH.OH, HC.H;0z, FeCls and NH.C2H:02.
Dilute and boil. (4)

Precipi- | Solution: Groups III, IV | Precipitate: | Filtrate: Nit+t, Co,++
tate: and V. Saturate with| FePO, and (Zn*+)Batt, §rtt,

(H2SnO;)x-| HS and filter. (2) Fe(OH).C2H;02.| Cott, Mgtt. Add
(P20s)y. . Reject. NH,OH- and saturate

Reject. with HS. (5)

(NH4)28S the filtrate

and con- obtained

| tinue as in after the

Table XI. jirst treat-

(3) ment with
NH,OH
and
(NH,4)8S
and exam-
ine for
Groups IV
and V as
in Table
XI. (6)

Basic AcETATE METHOD

4, Treat the filtrate from Group II with NH,OH and (NH,).S in the usual
way. The precipitate may contain FeS, NiS, CoS, ZnS, MnS, Al(OH) , AlPOx,,
Cr(OH)s, CrPOQ,, Bas(PO,)s, Srs(PO,)s, Ca;(POx)2, and MgNH,POs,. The
filtrate may contain Bat*, Srt*+, Cat+, Mgt+, Nat, Kt and NH,*. Set
the filtrate aside to be combined later with a solution which will contain
any Batt, Srt*, Cat* and Mg++ that was precipitated as phosphate by
NH.OH.

Dissolve the precipitate in HCl, adding a little HNO; if necessary. Dilute,
filter off any residual sulfur and evaporate the filtrate nearly to dryness to
remove the excess acid. Dilute to 20 cc., make alkaline with NaOH and add
more water if a very bulky precipitate is formed. Add about 2 cc. of Na,O0:

446 - SYSTEMATIC ANALYSIS

powder in small portions while stirring the cold solution. Then add 5 ce. of
4-normal Na2CQ; solution and boil to decompose the excess of NazOz. Cool,
dilute with an equal volume of water and filter. The filtrate will contain all
the aluminium as sodium aluminate, all the chromium as sodium chromate,
usually most of the zine as sodium zincate and some or all of the phosphoric
acid as sodium phosphate. Examine this filtrate for aluminium, chromium and
zinc exactly as outlined in Table VI, p. 192. The presence of the phosphate does
no harm, as zinc phosphate is readily soluble in ammoniacal solutions. When
the solution which was treated with sodium hydroxide and sodium peroxide
contains alkaline earths, the phosphates of these elements are precipitated by
the alkali and if the solution contains no alkaline earth metal, usually most of
the phosphate remains in solution. The sodium carbonate is added with the
sodium hydroxide and sodium peroxide to ensure the complete precipitation
of alkaline earth and to prevent the precipitation of an alkaline earth chromate.

Dissolve the precipitate produced by NaOH, Na,O. and NazCO; in 15 ce.
of 6-normal HNO;, adding HO. a few drops at a time till precipitate is all
dissolved. Filter if any filter paper is present and evaporate the filtrate nearly
to dryness. Add 16-normal HNO; and treat with KCIO; exactly as described
on p. 193. If a precipitate of MnO, is formed, filter it off and, if deemed
necessary, confirm the manganese in the usual way. Take one-tenth of the
filtrate from the KCIO; treatment, evaporate it just to dryness, moisten the
residue with a few drops of concentrated HCl, dilute with 5 cc. of water and
test for iron with KCNS. Neutralize the remainder of the solution with
ammonia and continue adding ammonia until a permanent precipitate is formed
or the solution becomes alkaline. Dissolve the precipitate in acetic acid,
avoiding an excess, and add 15 cc. of 3-normal NH,C,H;0, solution. If the
solution is then of a rich brownish red color, it is evident that more than
sufficient iron is present to combine with all the phosphoric acid. Otherwise
add FeCl; solution, drop by drop until such a color is produced. Dilute to
at least 100 cc. and hoil for five minutes in a 500-cc. flask. If a large precipitate
is formed, dilute to 250 cc. with hot water and boil a minute longer. Allow
the precipitate to settle for a minute or two, filter while still hot through a
plaited filter and wash the precipitate with hot water containing a little
ammonium acetate. Add 10 cc. more of NH,C.H;0; to the filtrate, again boil
and collect any further precipitate on a separate filter. Reject both of these
precipitates, which contain all the phosphoric acid as pale yellow ferric phos-
phate and the excess of the iron as basic ferric acetate.

5. Make the filtrate ammoniacal and saturate it with H.S. Filter and
examine the precipitate for nickel, cobalt and zinc according to the directions
on p. 193.

6. Add this last filtrate, which may contain Batt, Sr++, Cat* and Mgtt
as well as the added NH, salt, to the filtrate obtained by the original treat-
ment with NH,OH and (NH,).S. Examine the combined filtrates for Groups
IV and V as outlined in Table XI, p. 440.

EXAMINATION FOR THE NEGATIVE ELEMENTS (ANIONS)

The tests for the acids (anions) are usually made after the analy-
sis for the metals (cations); the preliminary examination (heating in
the closed tube and with dilute and concentrated sulfuric acid) and

_ EXAMINATION FOR THE NEGATIVE ELEMENTS 447

the solubility, combined with the knowledge of the metals that are
present, suffice to tell us what acids may and what acids may not be
present.

In order to avoid side-reactions, the acids are usually obtained in
the form of the neutral alkali salts before proceeding to test for
them.

PREPARATION OF THE SOLUTION FOR THE ANALYSIS FOR ACIDS

Two cases may be distinguished:

A. The original substance contains no heavy metal (i.e., only alkalies
or alkaline earths are present).

(a) The substance is soluble in water.

Test the solution with litmus paper to see whether it is acid, alkaline,
or neutral.

An Alkaline Reaction shows the possible presence of alkali cyanides, alkali
nitrites,* borates, tertiary phosphates, alkali sulfides, thio salts of the alkalies,
alkali silicates, etc.

An Acid Reaction is shown by many acid salts (cf. p. 49).

Divide the solution into two parts. If it is neutral, analyze it directly for
the acids; if it is alkaline, neutralize + half of it with acetic acid and the
other half with nitric acid; if it is acid, neutralize with sodium carbonate
solution.

(b) The substance is insoluble or very difficultly soluble in water, but
readily soluble in dilute acids. In this case only the acids of Groups III and
IV need be tested for.

Boil the dry substance with a little concentrated sodium carbonate solution
and filter. The filtrate contains the acids in the form of their sodium salts.

Neutralize the solution with dilute nitric acid.

The following substances may be present: BaSQ., SrSOu, (CaSO,), CaF»,
and silicates, which often contain salts of H;PO., HBO:, H.SO,, HF, and HCl.

Fuse the substance with sodium carbonate in a platinum crucible, extract
with water, and use the aqueous solution thus obtained for the analysis for
acids, after neutralizing.

If the substance is partly soluble in water and in acids, first treat with water
and then with sodium carbonate solution and fuse the residue with solid
sodium carbonate. Analyze separately the three solutions thus obtained. —

B. The substance contains heavy metals.

(a) It is Soluble in Water or in dilute acids, and contains no non-volatile
organic matter (no carbonization in the closed-tube test.)

Treat the solid substance with suificient concentrated sodium carbonate

* Perfectly pure alkali nitrites are neutral. The alkaline reaction of the com-
mercial salts is due to the presence of alkali oxide or silicate.

{ Taiosalts, silicates, stannites, stannates, aluminates, molybdates, tung-
states etc., will yield precipitates at this point which should be examined accord-
ing to Table XIII.

448 - §YSTEMATIC ANALYSIS -

solution to make the resulting solution weakly alkaline, and filter. If ammo- —
nium salts are present, first boil it with the solution of sodium carbonate until —
the vapors from the solution no longer smell of ammonia, and then filter. “all

Divide the resulting solution into two parts, making one part acid with —
acetic acid, and the other with nitric acid. a

(b) The Substance is Soluble in Water or Dilute Acids and Contaiuas |
Non-volatile Organic Matter.—If the metals of the ammonium sulfide and
hydrogen sulfide groups are both present, pass hydrogen sulfide into the
weakly acid solution until it is saturated, filter off the precipitate, add ammonia
to the filtrate until it is slightly alkaline, filter again, and make this last filtrate —
acid with acetic acid and evaporate to a small volume. Filter off the deposited
sulfur, treat the solution with solid potassium carbonate, filter if necessary,
carefully acidify with nitric acid, stir vigorously, and if any potassium-acid —
tartrate is formed, filter it off and test.as described on p. 363. Test the —
filtrate for the remaining acids. a

(c) The Substance is Insoluble in Strong Acids.—Besides the salts’ men-
tioned under A (c), the following may be present: AgCl, AgBr, AgIl, AgCN,
PbSO,, silicates (ferro- and ferricyanides).

If silver is present, the halogen acids must be looked for. Reduce the =
insoluble silver salt by zinc and sulfuric acid, filter off the residue, and examine -
the filtrate according to Table XV for HCl, HI, HBr, and HCN. ; 4

If the insoluble substance contains lead, boil it with sodium carbonate —
solution and filter; make the filtrate acid with hydrochloric acid, and test with _
BaCl. for H.SQ,. :

If silicates are present, H;PO., HF, HBO., HCl, and H.SO, must also be —
tested for. _

In whatever way a solution is prepared, determine its behavior toward |
silver nitrate and barium chloride in order to ascertain to what groups the ¥
acids present belong. 3

Then make the necessary tests for the individual members.

EXAMINATION FOR THE NEGATIVE ELEMENTS

449

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450

EXAMINATION FOR THE NEGATIVE ELEMENTS 451

TABLE XV.—EXAMINATION OF Group I

First test for HCN by placing a little of the solution on a watch-glass, adding a few
drops of yellow ammonium sulfide, evaporating carefully to dryness, acidifying
the dry mass with HCl, and adding a drop of FeCls solution. If a blood-red col-
oration is produced, HCN is present, in which case treat a larger portion of the
neutral solution with nickel sulfate * solution in excess, and filter.

PRECIPITATE. SOLUTION.
Ni(CN)2 Treat the solution, which is now free from hydrocyanic acid, with a
Discard. a little caustic soda solution (free from halogen), boil and filter off

the precipitate of Ni(OH)>2, divide the filtrate into two parts and use
one part for the HBr and HI tests and the other for the HCl test.

Tests for HI and HBr.

Test for HCl.

Make the solution acid with dilute

H.SO;, add chlorine water drop
by drop, and shake the solution
with CS: or CHCl;. If the latter
is colored violet, HI is present.
By further addition of chlorine
water, the CS, or CHCl; ts de-
colorized completely if HBr is
absent, but turned yellowish-
brown if HBr is present. If too
much chlorine water is used a
wine-yellow color is produced.

Make the solution slightly acid
with HNO;, and add dilute
AgNO; drop by drop. AglI
and AgBr are first precipitated
(yellow). Filter and add more
AgNO;. If the precipitate
still appears yellow, filter
through a new filter, and
again add AgNO; to the filtrate
until a white precipitate of
AgCl is formed if HCl is
present.

* If ferricyanic acid is present also add a little ferrous sulfate.

precipitated by nickel sulfate.

Group II

Ferrocyanic acid is completely

The members of this group are almost always detected in the preliminary
examination. The special tests for these acids are described on p. 329 et seq.

Group III

SO2, COz, H2C.0, are recognized in the preliminary examination. HPOs,
H,P,0,, HBO,, and H.C,H,O, are tested for separately by the special reactions
described on p. 347 et seq.

Group IV

Cr0;, H;PO,, and H.8.0; are detected in the preliminary examination, and
in the analysis for metals.

Group V

HClO; and HNO; are usually detected in the preliminary examination.
Their presence is, however, always confirmed by the procedure described on

p. 398.

Groups VI and VII

These acids are usually detected in the preliminary examination. Their
presence is confirmed by the tests described under H.SO,, HF, and SiOx.

452 SYSTEMATIC ANALYSIS

B. THE SUBSTANCE IS A METAL OR AN ALLOY

The analysis of a metallic alloy is much simpler than that of a mixture of
salts, because there are no acids to test for. Of the electro-negative elements,
usually only phosphorus, silicon, carbon, and sulfur have to be considered.

As all metals, with the exception of gold, platinum, tin, and antimony, are
soluble in nitric acid, alloys are usually brought into solution by dissolving
therein, and the use of aqua regia is only necessary in a few cases. Many

alloys rich in silicon (e.g., copper silicide) are extremely difficultly soluble even 4

in aqua regia, and are best brought into solution by fusing with caustic alkali.

in a silver crucible, and afterwards dissolving the melt in nitric acid.

It is not advisable to dissolve an alloy in hydrochloric acid, for phosphides,
carbides, silicides, sulfides, and arsenides, which are often present in small
amounts, are decomposed by this acid in such a way that the negative elements
are evolved as hydrogen compounds, and thus escape detection. For the
analysis of ordinary alloys, the following procedure is used:

Place 1 or 2 gms. of the alloy (best in the form of borings) in a 200-ce. porce-
lain dish and treat them under a good hood ‘with about 20 ce. of nitric acid, sp. gr.
1.25-1.30 (1 vol. cone. HNO;+1 vol. H:O). After the first violent reaction

is over, carefully evaporate (with constant stirring) almost to dryness, being —

careful to avoid overheating; * add a little water and heat.

(a) The mass dissolves completely. The alloy contains neither tin nor
antimony; analyze it according to Table XVI.

(b) The mass does not dissolve completely, but a white, greenish residue
remains; analyze according to Table XVII.

C. THE SUBSTANCE IS A LIQUID

The color, odor, and reaction towards litmus enable one to draw important
conclusions.

(a) The solution reacts neutral; it contains no free acid, free base, acid

salt, no salt which shows an acid or alkaline reaction on account of hydrolysis,
nor any insoluble salt.

First of all, determine whether there are any solid substances dissolved in —

the liquid by evaporating a small portion to dryness at as low a temperature
as possible (so as not to lose any volatile substances). If a residue remains,
examine it according to A, p. 428.

(b) The solution reacts alkaline. An alkaline reaction may be due to the

presence of hydroxides of the alkalies or alkaline earths, peroxides, carbonates,
borates, cyanides, silicates, sulfides (zincates, aluminates, molybdates, tung-
states) of the alkalies, as well as by ammonia or hypochlorites, ete.

If the solution, for example, contains hydroxides or carbonates of the alkalies,
it is evident. that substances which are precipitated by them cannot be present
at the same time, except, in some cases, in the form of complex ions (cyanides,
tartrates, etc.)

At once test the solution for peroxides, hydroxides, and carbonates, as well
as for the sulfides of the alkalies.

* Otherwise insoluble basic salts are likely to be formed. If this be the case, —

as is often shown by the dark color of the residue, add a little conc. HNOs, heat
the liquid somewhat and then dilute with water.

453

THE SUBSTANCE IS A METAL OR AN ALLOY

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SYSTEMATIC ANALYSIS

454

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}

THE SUBSTANCE IS A METAL OR AN ALLOY 455

To test for peroxides * (H.0.) heat a little of the solution with a few drops
of cobalt nitrate solution; a black precipitate shows the presence of H2O:.t
Or test the solution by adding some titanium sulfate solution and acidifying
carefully with cold dilute sulfuric acid; a yellow coloration shows the presence
presence of H,0..t

A still more sensitive reagent, according to Schéne,§ is a very dilute solu-
tion of FeCl;+K;[Fe(CN),]. If the slightest trace of H.O,. is present in the
solution, the red solution becomes greenish, and after a time Prussian blue
separates out.

In order to detect the presence of hydroxides and carbonates in the presence
of H;02, boil a portion of the solution for a long time in a porcelain dish in order
- to destroy the peroxide, and then add barium chloride until no more precipitate is
formed. If the solution now shows an alkaline reaction, the presence of hydrox-
ides || is assured. If the precipitate produced by barium chloride dissolves
in acid with effervescence, and the gas evolved renders barium hydroxide
solution turbid, carbonates are present. If the solution smells of ammonia,
evaporate a small portion to dryness in order to see whether other compounds
are present, and examine the residue according to A, p. 423.

(c) The solution reacts acid; it can then contain substances which are
soluble in water and in acids, as well as free acids. Evaporate a small portion
to dryness in order to see whether any non-volatile matter is present. If no
residue is obtained, neutralize the solution with soda and test for acids. If
a residue is obtained, examine it according to A, p. 423.

D. The substance to be analyzed is a gas.

This case will be considered in Volume II under Gas Analysis.

* See foot-note, p. 428.

{If the alkaline solution contains hydrochlorites or sulfides, these will also
give a black precipitate with cobalt nitrate; the above reaction serves only to detect
HO, in the absence of hypochlorites or sulfides. If the latter bodies are present,
H,O2 cannot be present, because hypochlorites are reduced to chlorides and sul-
fides oxidized to sulfates by H2O.. f

The presence of hypochlorites is usually detected by. the odor, on acidifying
with dilute H.SO,, the odor of chlorine can be detected. Sulfides give off H.S on
being acidified. Hypochlorites and su'fides cannot exist together in the same
solution. .

t H.O2 can also be detected by the chromic acid reaction, but this test is less
certain than that with titanium sulfate. |

§ Ber., 7, 1695.

|| Either as such in the original solution or by hydrolysis of peroxides. This
test for OH ions in the presence of carbonates has been used for years by the
author. It is reliable, though care must be taken to add an excess of BaCh, as
otherwise an alkaline reaction may be due to BaCO;, which is more soluble in
water than in BaCl, solution.

PART V.—REACTIONS OF SOME OF THE RARER
METALS

In the treatment of the rare metals, the same method and order
will be used as with common metals.
THE ALKALI GROUP
CASIUM, RUBIDIUM, LITHIUM
CESIUM, Cs. At. Wt. 132.81. M. Pt.=26°

Occurrence.—Cesium and rubidium are not rare elements, strictly.

speaking, for they are found almost everywhere, but always only

in very small amounts. Thus cesium replaces potassium in many
feldspars and micas, and is found in many rocks which carry these
minerals, as well as in mineral waters which ooze from them. Czsium _
and rubidium were discovered in the mother liquor of Durkheimer
brine in the year 1860 by Bunsen and Kirchhoff by means of the spec
troscope. ae

Pollucite, a mineral closely related to leucite, found at Elba and ~
crystallizing in the regular system, is a typical cesium mineral. Its
composition is H2@s4Al4(Si03)9.

Cesium and rubidium in all their reactions behave almost exactly
like potassium. The principal difference between these three metals

lies in the different solubilities of their corresponding salts, as will @

be seen by the table given on p. 459.

REACTIONS IN THE WET WAY

A solution of cesium chloride should be used. }

1. H2[PtCls] produces a yellow, crystalline precipitate which is —
of a lighter color than the corresponding potassium salt and is much —
less soluble; 100 parts of water dissolve at 0° C. only 0.024 part and at
100° C. 0.377 part of the salt.

2. Tartaric Acid produces, as with potassium and rubidium, a a

white, crystalline precipitate, CsHC4H40.; 100 parts of water dis- — |
solve at 25° C. 9.7 parts, and at 100° 97.1 parts of the salt. 9

3. He[SnCl¢] (a solution of SnCl, in concentrated HCl) produces —

in concentrated solutions a white, crystalline precipitate of Cs2{SnCle] q
456

THE ALKALI GROUP 457

(octahedrons). Ammonium salts give the same reaction, but potas-
sium and rubidium salts do not.

REACTIONS IN THE DRY WAY

Cesium compounds color the flame-reddish violet, very similar
to the potassium flame.

Flame Spectrum. An intensely blue double line, 455.5upu, 459.3up.
At higher temperatures a number of paler lines appear of subordinate
importance; in the red 697.3uu and 672.3uy, in the orange yellow
621.3uu and 601yup, in the yellow 584.5yy, in the green 566.4uu, 563.5up,
550.3uu, 547.1uy, 541.9un, and 535.luu. Moreover, a faint. contin-
uous spectrum is seen from the yellow to the blue. See chart.

RUBIDIUM, Rb. At. Wt. 85.45. M. Pt.=38°

Occurrence.—Rubidium almost always accompanies cesium and is
found in many mineral waters; in carnallite from Stassfurt; in lepido-
lite, (Li,K,Na)2Ale(F,OH)2Sis09; in triphylite, (FeMn)(LiCsRb)PO,;
and in spodumene, (LiNa)AIl(SiO3)2 a mineral of the pyroxene group.
Lepidolite from Rozena contains about 0.54 per cent Rb and 0.0014
per cent Cs. As far as the author knows, there is no typical rubidium
mineral.

REACTIONS IN THE WET WAY

1. H»[PtCls] produces, as with cesium and potassium salts, a
white, crystalline precipitate of Rbe[PtCle], which is more difficultly
soluble than the corresponding potassium salt, but more soluble
than the cesium salt; 100 parts of water dissolve at 0° C. 0.134 part,
and at 100° 0.634 part of the salt.

2. H2[SnCle] produces a. white precipitate only in very concen-
trated solutions. The salt is much more soluble than the corre-
sponding cesium salt, but this reaction is not suitable for separating
the two metals.

3. Tartaric Acid produces a precipitate of RbHC.zH40¢ only in
concentrated solutions; 100 parts of water dissolve at 25° C. 1.18
parts, and at 100° C. 94.1 parts of the salt. The corresponding cesium
salt is more soluble, while that of potassium is less soluble.

REACTIONS IN THE DRY WAY

Flame Coloration.—Similar to csesium.
Flame Spectrwm.—Violet double lines 420.2uu and 421. Bias also
the red double line 781.1uu and 795.0uy. At higher temperatures a

458 REACTIONS OF SOME OF THE RARER METALS

continuous spectrum is visible from the yellow to the blue in which
the following lines are to be found in addition to those given above;
in the orange-yellow 62).8uu, 626.1uu, 620.64, 617.7up; in the yellow-
green 572.4uu, 570.0uu, 564.8un and of subordinate importance the
ereen lines 543.5upu, 536.5uu, and 516.8up.

LITHIUM, Li. At. Wt. 6.94. M. Pt.=186°

Occurrence.—Lithium is found to a greater extent in nature than
cesium and rubidium; in triphylite, (Fe,Mn)(Li,Cs,Rb)PO4; in
petalite, Al(Li,Na,H)Si40i0, a mineral of the feldspar group (also
called castorite); in amblygonite, Li(AIF)PO4, monoclinic; in lepido-
lite, Ale(Li,K,Na)o(F,OH)2Siz09; in many varieties of tourmaline,
and muscovite, in epidote and orthoclase, and consequently in many

mineral-spring waters. In some cases as much as 36 mgs. Li are

contained in a liter of spring water.

Lithium is the lightest of all metals, and floats on petroleum.
It oxidizes quickly in the air, and decomposes water at ordinary tem-
peratures, forming LiOH, which dissolves slowly in the water; the
solution reacts alkaline and absorbs carbon dioxide from the air
with avidity, forming difficultly soluble LigCOs3.

Lithium chloride is soluble, even in the anhydrous state, in a mix-—

ture of alcohol and ether as well as in amyl alcohol (difference from
the remaining metals of this group).

REACTIONS IN THE WET WAY

A solution of lithium chloride should be used.

1. He{PtCle] produces no precipitation.

2. Tartaric Acid produces no precipitation.

3. NazHPO, produces from boiling, moderately concentrated solu-
tions a white precipitate of tertiary lithium phosphate. ‘lhe pre-
cipitation is only quantitative when the solution is made alkaline
with caustic soda, evaporated to dryness, and taken up in water con-
taining ammonia:

HPO4+3Lit+OH~ — LisPO4+ H20.

Lithium phosphate is fusible (difference from magnesium and
the alkaline earths). .
4. (NH4)2CO3. If ammonium carbonate and ammonia are added

to. a concentrated lithium solution, lithium carbonate is precipitated

in the form of a white powder. The salt, contrary to the other alkali

THE ALKALI GROUP 459
carbonates, is: very difficultly soiuble in water: 100 parts of water
dissolve at 13° C. 1.31 parts of LigCO3. In the presence of con-
siderable alkali chloride or of ammonium chloride, no precipitation
' takes place.

REACTIONS IN THE DRY WAY

Flame Coloration.—Pure lithium salts impart a magnificent, car-
mine-red coloration to the gas-flame. If considerable amounts of
sodium salts are present at the same time, the lithium flame is com-
pletely masked; but if the flame is observed through cobalt glass the
red color becomes distinctly visible.

Flame Spectrum.—An intensely red line 670.8up and at higher
temperatures the pale orange-yellow line 610.8uy.

SUMMARY OF THE ALKALI METALS

LITHIUM. Sopium. Porassium.} Ruprprium.| Casrium.
Atomic WEIGHT..... 6.94 23.00 39.10 85.45 132.81
MELTING-POINT....... 180° 95.6° 62.5° 38 .5° 26-27°
X.[PtCl,]:
Solubility in alcohol soluble soluble insoluble | insoluble | insoluble
odes sp aes considerable | considerable} 1.12 0.141 0. 079
ah 100s. s considerable | considerable | 5.13 0.634 0.377
XHC,H.Os:
Pe eae 1 i sat considerable | considerable | 0.425 — a
1 he? lal ea _ — — 1.18 9.7
ALUMS:
aE er bul a \ ne = 13.5 2.27 0.619
CHLORIDES: ‘
Sapte = oa soluble insoluble | insoluble | insoluble | insoluble
CARBONATES: :
age eee ome insoluble insoluble | insoluble | insoluble | soluble

DETECTION OF LITHIUM, RUBIDIUM, AND CASIUM
IN THE PRESENCE OF CONSIDERABLE SopIuM AND POoTAssIUM

Evaporate the solution containing the metals as chlorides almost to dryness,
triturate the residue with 90 per cent alcohol and filter. The alcoholic solution
contains all of the Li, Rb, and Cs, with small amounts of K and Na. Evaporate
again nearly to dryness and once more extract with alcohol and filter. If only
traces of the rare alkalies are present, the residue must be extracted several
times with alcohol.

Evaporate the alcoholic extract to dryness and treat the residue with con-

460 REACTIONS OF SOME OF THE RARER METALS

centrated hydrochloric acid. This is done in order to change LiOH into LiCl. —
Some of the former is formed by evaporating the solution, and it is insoluble —
in alcohol-ether. Evaporate once more, gently ignite the residue over the free.
flame, and, after cooling, triturate the residue with ether-aleohol mixture, using
a elass rod, and filter through a filter that is wet with the ether-alcohol mixtures .
The filtrate contains the lithium. Evaporate it to dryness and test for lithium | 4
by means of the flame reaction or in the spectroscope. a

Dissolve the residue insoluble in alcohol-ether in a little water, treat with —
chloroplatinic acid and filter. Extract the precipitate repeatedly with small —
portions of boiling water, decanting off the liquid each time through the filter. |

The potassium chloroplatinate dissolves in the hot water, forming a —
yellow solution. Continue the treatment with hot water until the residue is —
of a light-yellow color. Dry, place in a porcelain boat, and heat in a glass —
tube, made of difficultly fusible glass, in a stream of dry hydrogen; in this —
way the alkali chloroplatinates are reduced to chloride and platinum:

X{[PtCh]-+2H2 =4HCl ¢ +2XC1+Pt.

After cooling, treat the residue with a little water, filter off the platinum, —
evaporate the solution to dryness and test in the spectroscope for: Cs and Rb. |

In order to detect lithium, cesium, and rubidium in a silicate undecom-—
posable by acids (lepidolite, for example) decompose the finely powdered —
mineral with hydrofluoric and sulfuric acids as described on p. 419, change
the sulfates to chlorides by the addition of barium chloride, and ‘free the >
solution from other metals as described on p. 96, and then carry out the
above separation. |

METALS OF THE (NH,),S GROUP

BERYLLIUM, ZIRCONIUM, THORIUM, YTTRIUM, ERBIUM, CERIUM,
LANTHANUM, DIDYMIUM, TANTALUM, NIOBIUM

BERYLLIUM, Be. At. Wt. 9.1. M. Pt.=>1880°

Occurrence.—Chrysoberyl, Be(AlOz)2; phenacite, Be2SiO4; _bery],
BezAleSigO1g; euclase, AlBeHSiOs; meliphanite, Be2CazNaSizO10F;
and leucophanite, BeCaNaSi2OcF.

Beryllium is a bivalent metal, and forms a white oxide, BeO,
which is soluble in acids. Beryllium salts react acid in aqueous solu-
tion and possess a sweetish, astringent taste. The element is often
called Glucinum, Gl, in England and the United States.

REACTIONS IN THE WET WAY

Use a solution of BeSO4-4H20.

1. Ammonia and Ammonium Sulfide produce a white precipitate
of Be(OH)2, similar in appearance to AlfOH)s, insoluble in an excess
of the precipitant, but readily soluble in HCl, forming a colorless
solution. The yellow color often obtained in dissolving the hydroxide
in hydrochloric acid is due to traces of ferric chloride.

2. KOH precipitates white, gelatinous beryllium hydroxide, readily
soluble in an excess of the reagent, forming potassium beryllate:

Be**++20H- — Be(OH)2; Be(OH)2+20H~ — [BeO2]~+2H20.

The alkali beryllates are decomposed hydrolytically on boiling their dilute
aqueous solutions, all of the beryllium being precipitated as hydroxide. The
precipitate thus obtained is denser than that thrown down by ammonia,
and differs from the latter by being insoluble in potassium carbonate and
difficultly soluble in ammonium carbonate solutions; it is also much more
difficultly soluble in dilute acids. A solution containing a considerable excess
of alkali hydroxide does not give a precipitate of beryllium hydroxide by
boiling.

3. Ammonium Carbonate produces a white precipitate of beryl-
lium carbonate, readily soluble in an excess of the reagent (difference
from aluminium); by boiling the solution, the beryllium is precipi-
tated as white basic carbonate. This property enables us to sepa-
rate beryllium from iron and aluminium. The separation is not

461

462 REACTIONS OF SOME OF THE RARER METALS

sharp, however; which can also be said of the separation by means of
caustic potash.

In order to make a quantitative separation, the beryllium hydroxide or 7g

carbonate must be redissolved and the precipitation repeated several times.

4. BaCO3 precipitates beryllium completely in the cold as hydroxide. —

5. Oxalic Acid and Ammonium Oxalate cause no precipitation
(difference from thorium, zirconium, erbium, yttrium, cerium, lan-
thanum, and didymium).

6. KeSO4 gives with beryllium sulfate a beautifully crystalline
double salt, Ke2Be(SO4)2-2H20, which is soluble in a concentrated
solution of KeSOx4 (difference from Ce, La, and Di). |

7. BeCle is soluble in a mixture of equal volumes of saturated
aqueous and ethereal hydrochloric acid, while the hydrous alumin-
ium chloride is not (good method for separating Be and Al).*

There are no characteristic dry reactions for beryllium.

ZIRCONIUM, Zr. At. Wt. 90.6. M.Pt.<Si

Occurrence.—Zircon, Zr28iO4, tetragonal, isomorphous with rutile
TizO4, thorite (orangite), ThSO., cassiterite, Sn2O4, polianite,
Mn204, and plattnerite, Pb2O4; baddeleyite, ZrO2, monoclinic.

Zirconium forms two oxides: Zirconium dioxide, ZrOez, and azir-
conium pentoxide, Zr2O5. The former is the more important and can
be dissolved by heating for a long time with a mixture of two parts
of concentrated H2SO4 and one part of water and afterwards diluting.

The mineral zircon, ZrSiO., cannot be decomposed by such treatment. It
must be finely pulverized and fused with four times as much sodium carbonate
at a high heat in a platinum crucible; sodium silicate, Na,SiO., and sodium
zirconate, NasZrQO,, are formed. On treating the melt with water, the former
salt dissolves, while the latter is decomposed hydrolytically, forming sodium
hydroxide and sandy, insoluble zirconium hydroxide; the latter retains some
of the caustic soda very persistently. After washing the residue, heat it, without
previous drying, with concentrated sulfuric acid at a temperature near
boiling-point; in this way anhydrous Zr(SO,)2 is obtained. By pouring
water over the latter, the salt Zr(SO,)2-4H2O0 is formed, which dissolves slowly
in cold water, but more readily in hot water, forming a solution with an acid
reaction.

REACTIONS IN THE WET WAY
A solution of zirconium nitrate, or a freshly prepared one of are
oxychloride, may be used for the following reactions:

1. NH4OH and (NH4)2S produce a white gelatinous precipitate
of Zr(OH)a4, insoluble in an excess of reagent.

*F.S. Havens, Z. anorg. Chem., 18 (1898), 147.

;
‘
!
:
|

METALS OF THE (NH,).S GROUP 463

2. KOH and NaOH likewise produce the same precipitate insol-
uble in an excess of reagent (difference from Al and Be). When the
zirconium hydroxide is produced in the cold it is readily soluble in
dilute acids; but when thrown down from a boiling solution it is very
difficultly soluble in dilute acids, though it will dissolve even then in
concentrated acids without difficulty.

3. (NH1)2CO3 produces a white, flocculent precipitate of basic
carbonate, readily soluble in an excess of the reagent, but reprecipi-
tated by boiling.

4. K2COs and NazCOz produce white precipitates somewhat
soluble in an excess, but reprecipitated by ammonia.

5. BaCO3 causes incomplete precipitation, even on boiling.

6. Oxalic Acid gives a white, ftocculent precipitate of zirconium
oxalate, readily soluble in an excess of oxalic acid, difficultly soluble
in dilute hydrochloric acid, and readily soluble in ammonium oxa-
late. From the solution in (NH4)2C2O4 the zirconium is not pre-
cipitated by the addition of dilute HCl (difference from Th).

7. Ammonium Oxalate behaves the same as oxalic acid.

From the solution in ammonium oxalate, zirconium is not pre-
cipitated on the addition of hydrochloric acid (difference from tho-
rium).

Remark.—A solution of zirconium sulfate behaves quite differently from
that of the nitrate and oxychloride towards oxalic acid and ammonium oxalate,
a fact which, although published by Berzelius and also by Pfaff, had been
entirely forgotten by most chemists until their attention was called to it by
R. Ruer.*

On treating an aqueous solution of zirconium sulfate with oxalic acid or
ammonium oxalate there is no precipitation; in fact, precipitation will not
take place from nitrate or chloride solutions when these contain sufficient
sulfuric acid, sodium or ammonium sulfate.

The cause of this different behavior lies in the fact that zirconium forms
complex compounds with sulfuric acid and alkali sulfates. Thus the solution
of zirconium sulfate contains the acid H,[ZrO(SO,).], and on treating a solution
of the oxychloride or nitrate with sodium or ammonium sulfate (but not the
potassium salt) the sodium or ammonium salt of this complex acid is formed:

ZrOCl,+2Na80, =2NaCl+ Na,[ZrO(SO,),].

These compounds, however, are electrolytically dissociated in aqueous
solution as follows:

H,[ZrO(SO,)] = 2H*++ [ZrO(SO,).]*.
As the zirconium is present in the anion it ‘erint react with oxalic acid.
8. HF causes no precipitation (difference from Th and Y).
* Z. anorg. Chem., 42, 85 (1904).

464 REACTIONS OF SOME OF THE RARER METALS

9. K2SO4.--A concentrated, cold solution of K2SO4 precipitates —
little by little all of the zirconium as potassium zirconium sulfate, —
insoluble in an excess of the reagent (difference from Al and Be). —
The precipitate, when produced in the cold, dissolves readily in con-—
siderable dilute HCl. If it is produced from a boiling solution, basic —
zirconium sulfate is formed by hydrolysis, which is quite insoluble —
in dilute HCl (difference from Th and Ce). oe

10. NasSO4 produces no precipitation, even on boiling the solu-
tion, which is slightly acid with sulfuric acid (difference from Ti). —

11. H2O2 precipitates from slightly acid solutions white; volu-
minous zirconium peroxide, Zr2Os5, which evolves chlorine on being —
warmed with concentrated HCl. |

12. NaeS2Oz precipitates zirconium completely as the hydroxide, a
the precipitate being always contaminated with sulfur. |

13. Turmeric Paper, after being moistened with the hydrochloric —
acid solution of a zirconium salt and dried, is colored reddish-brown
(difference from Th). a

14. HCl. Ruer * recommends the following test for the identi- —
cation of zirconium:

Precipitate the zirconium in the cold by ammonia, filter, wash, and separate
it from the filter as completely as possible. Dissolve the precipitate in hydro- —
chloric acid (or if small in amount treat the paper and precipitate together —
with not too strong HCI and filter.) Evaporate the hydrochloric acid solution —
to dryness on the water-bath and take up the residue in a little water. To the —
cold, saturated solution add hydrochloric acid drop by drop, when the presence —
of zirconium will be evident by the formation of a voluminous precipitate of —
zirconium oxychloride. Redissolve the precipitate by heating the solution, ~
and allow the liquid to cool. After some time fine, silky needles of ZrOCl,.8H,O —
will precipitate.

In the somewhat unusual case that zirconium is present in the form of the 4
insoluble meta-zirconium acid, transform the latter into zirconium sulfate by 3
heating with concentrated sulfuric acid (2:1), dissolve this in water, precipi- —
tate the zirconium by ammonia, and carry out the above process. 3

15. Sodium Iodate produces in slightly acid solutions a volumi- —
nous white precipitate of zirconium iodate which is soluble in hot, —
dilute hydrochloric acid (best method of separating zirconium from
aluminium). } .

16. Hydrofluoric Acid as a rule spodices no precipitation (differ- ~
ence from thorium, cerium and other rare earths). From concentrated
zirconium solutions a voluminous precipitate may be obtained by the —
careful addition of hydrofluoric acid, but the precipitate is soluble —
in an excess of the reagent. |

* Z. anorg. Chem., 42, 85 (1904).

METALS OF THE (NH,).8 GROUP 465

REACTIONS IN THE DRY WAY

ZrOs is infusible in the oxyhydrogen flame (difference from the
other earths), but glows brightly.

THORIUM, Th. At. Wt. 232.4. M. P.>1700<Pt.

Occurrence.—Thorite (orangite), ThSiO4, with 50 to 58 per
cent ThO2; thorianite, a mineral discovered in Ceylon, with 72
to 76 per cent- ThOg and 11 to 12 per cent UO2;* gadolinite,
Be(Y,Ce,La,Di,Th,O)2FeSiO4, monazite, (Ce,La,Di,Th)PO4, with 2 to
8 per cent ThO2; and in the rare niobates, samarskite, pyrochlore,
euxenite, etc. Euxenite is essentially a titanate and niobite of
Ce(La,Di) and usually contains UOz and FeO. Thorite, monazite,
and gadolinite are decomposed by acids, preferably sulfuric acid.

REACTIONS IN THE WET WAY |

A solution of Th(SO4)2 should be used.

1. (NH4)OH, (NH4)2S, or KOH produces a white precipitate of
Th(OH).4, insoluble in an excess of the reagent, but readily soluble
in dilute acids. By igniting the hydroxide, ThOz is obtained, which
is soluble in concentrated sulfuric acid only after long digestion.

2. KeCOz3 or NazCO3 precipitates the white carbonate, soluble in
an excess of the reagent, and not reprecipitated by the addition of
ammonia. On boiling, the solution becomes turbid, but clears again
on cooling. |

3. (NH2)COxz precipitates the white carbonate, readily soluble in
an excess; on warming to 50° a basic carbonate is precipitated,
which redissolves on cooling the solution. Ammonia causes no pre-
cipitation in this solution.

4. BaCOs completely precipitates thorium salts in the cold.

5. KeSOx4 precipitates K4Th(SOs4)4+2H20, difficultly soluble in
water and insoluble in concentrated K2SO4 solution (difference from
Y). The corresponding sodium compound is readily soluble in water.

6. Oxalic Acid precipitates, from solutions which are not too acid,
all of the thorium as white, crystalline oxalate, practically insoluble
in oxalic and dilute mineral acids.

7. Ammonium Oxalate likewise precipitates thorium oxalate,
which dissolves on boiling with a large excess of this reagent. The
solution remains clear after cooling, provided the original solution
did not contain too much free sulfuric acid, and enough ammonium

* Chem.-Ztg. Rep., 1905, 91.

466 REACTIONS OF SOME OF THE RARER METALS

oxalate was used. From the boiling solution of the ammonium
double oxalate, HCl precipitates practically all of the thorivm as
oxalate (difference from Zr).

In the presence of ammonium acetate, ammonium oxalate pro-

duces no precipitation; but by the addition of HCl almost all of the
thorium will be precipitated as oxalate.

8. HF produces a white, gelatinous precipitate, which soon changes
to a heavy powder. KF causes the same reaction.

9. Na2S2Oz3 precipitates all of the thorium on boiling.

There are no characteristic dry reactions.

SE ee eT ee

ee ee

”
es eo

_ 7, ‘
ve ae ll

THE GADOLINITE METALS

YTTRIUM, Y (At. Wt. 89). M.P. between 1000°-1400°, and ERBIUM,*
Er (At. Wt. 167.7)

Occurrence.—Yttrium is an important’ constituent of gadolinite,
Be(Y,Ce,La,Di,Th,Er,O)2FeSiO4, and of yttrotantalite Y(Nb,Ta)Os,
an isomorphous mixture of yttrium tantalate and yttrium niobate.
The two elements Y and Er are also found in cerite, thorite, and
monazite.

REACTIONS IN THE WET WAY
A solution of Y(NO3)3 and one of Er(NOs)3 should be used.
Yttrium Erbium

1. NH,OH and (NH4)2S pre- Behaves like yttrium.
cipitate the white hydroxide, in-
soluble in an excess.

2. KOH and NaOH precipi- Behaves like yttrium.
tate the white hydroxide, insoluble
in an excess; the presence of tar-
taric acid does not prevent precipi-
tation; but in this case yttrium
tartrate is precipitated (difference
from Al, Be, Th, and Zr). On
igniting the precipitate, the oxide
is obtained, which is readily soluble
in acids.

3. (NH4)2COs produces a white Behaves like yttrium, except
precipitate of the carbonate, read- | that the solution does not be-
ily soluble in an excess of the re- | come turbid on standing.
agent; after standing some time
the solution becomes turbid, owing

to the deposition of a double salt,
Yo2(COs)3 ° 2(NH4) 2CO3 -2H20.

*Erbium is not an element itself, but consists of at least three elements—
holium, thulium, and dysprosium. The separation of these elements is exceedingly
difficult; for this reason we shall consider simply the reactions of the mixture.

467

468
Yttrium

4. K2CO3 and NasCOs pre-
cipitate the white carbonate, read-
ily soluble in excess; after standing
some time an insoluble double salt
separates out.

5. BaCOgz does not precipitate
yttrium in the cold, and only in-
completely on warming.

6. Oxalic acid precipitates
white yttrium oxalate, insoluble
in an excess, difficultly soluble in
HCl, and perceptibly soluble in
ammonium oxalate.

7. K2SO. forms a double salt
which is soluble in KeSO,4 solu-
tion (difference from Zr, Th, Ce,
La, and Di).

8. HF produces white amor-
phous YF3, which becomes pulver-
ulent on warming, and is insoluble
in water and in HF (difference
from Al, Be, Ur and Ti).

Yttrium solutions do not give
an absorption spectrum.

REACTIONS IN

Yttrium oxide is strongly lumi-
nous on being heated; otherwise
there is no reaction.

REACTIONS OF SOME OF THE RARER METALS

Erbium

Behaves like yttrium, only the 4
solution does not become turbid.
on standing.

Erbium is not precipitated
at all, even on warming.

In erbium solutions, oxalic
acid produces a light-red, pul-
verulent precipitate; otherwise
the reaction is the same as with
yttrium.

Behaves like yttrium.

Erbium solutions give a char- |
acteristic absorption spectrum:.

Mu aye bp Me
683.9 640.5 523.2 468.5
667.1 548.9 491.6 449.9
653.5 541.0 487.5 422.3

649.0 536.4 474.5 416.6

THE DRY WAY

Erbium oxide, on being heated
on a platinum wire, colors the
flame distinctly green. If the
light is viewed through a spectro-
scope, a number of bright lines —
will be seen in the dark green
which coincide with the dark
lines obtained in the absorption —
spectrum. a

THE CERITE METALS

CERIUM, Ce. At. Wt. 140.6, M. Pt. 635°? LANTHANUM, La, At. Wt.
139.0. M. Pt. 810°

PRASEODYMIUM, Pr. At. Wt. 140.6. M. Pt. 940°?

sige an Secale Nd. At. Wt. 144.3. M. Pt. 840°?

Occurrence.—These three metals are important constituents of
cerite, H3Ca(Ce,Al)38ig013), and of orthite (allanite), HCazCes3‘ i3013,
besides being usually found associated with’ the gadolinite earths in
gadolinite, ete.

CERIUM

Cerium forms two oxides, Ce2Os and CeQs ; both are basic an-
hydrides, from which salts are derived. The cerous salts are white,
while the ceric salts are orange red.

REACTIONS IN THE WET WAY
1. Cerous Salts

A solution of cerous nitrate, Ce(NO;);, should be used.

1. NHsOH and (NH4)2S each produce a white precipitate of
Ce(OH)s3, insoluble in an excess of the reagent, but readily soluble
in acids. In the presence of tartaric and citric acids, etc., the above
reagents cause no precipitation (difference from Y). :

2. NaOH or KOH also precipitate white Ce(OH)s, even in the
presence of tartaric acid, ete. The white Ce(OH)3 becomes yellow
gradually on standing in the air, on account of being oxidized to
Ce(OH)a. 3

3. KeCO3 and (NH4)2COs each produce a white precipitate insol-
uble in an excess of the reagent.

4. Oxalic Acid or Ammonium Oxalate precipitate white cerous
oxalate, insoluble in an excess of the reagent, and in dilute mineral
acids. On ignition, pale yellow, insoluble CeOzg is formed.* If the
oxalate is contaminated with praseodymium oxalate, a cinnamon-
colored oxide is obtained, which is perfectly soluble in dilute acids.

* Only when the cerous oxalate is pure. If it contains traces of praseodymium,
the CeO, is obtained as a bright-yellow powder.
469

470 REACTIONS OF SOME OF THE RARER METALS

5. BaCOs slowly precipitates all of the cerium in the cold.

6. K,SO4. If a neutral solution of cerous salt is treated with
solid K2SO4 until no more will dissolve, a crystalline precipitate of
Ce2(SO4)3-3K2SO4 is slowly formed at room temperature, or more

quickly on heating. All of the cerium can be precipitated in this way,

as the double sulfate is insoluble in concentrated potassium sulfate
solution, but it will dissolve in considerable pure water or more readily
in acids. From slightly acid solutions cerium can be completely
precipitated with KeSO4 as Ce2(SO4)3-2K2S04-2H20. (Difference
from Al, Be and yttrium earths.)

Na2SOx4 behaves similarly (difference from Th and Zr).

7. HF produces in neutral and slightly acid solutions of cerous
salts a gelatinous precipitate which by long digestion with the hot
solution gradually becomes pulverulent. The precipitate of CeF3-H2O
is practically insoluble in water and dilute hydrofluoric acid, but readily
soluble in other mineral acids (difference from Al, Be, Zr and Ti.)

8. H2O2 colors neutral cerous solutions yellow, but after adding

an acid the’ color disappears, as the cerium peroxide is reduced to
cerous salt by hydrogen peroxide and acid. If a cerous salt is treated
with a slight excess of ammonium hydroxide and then with He2Oo,

the precipitate becomes reddish orange in color, something like Fe(OH)3.
_ This is the most sensitive test for cerium; it was discovered by Lecoq

de Boisbaudran. The composition of the precipitate has been given

as CeOo-Ce203-HeO2 and as Ce(OH)302H.

9. Chlorine. If a cerous salt is treated with an excess of alkali —

hydroxide and then with chlorine gas, a yellow precipitate of
CeO2-3H20 is obtained. If the chlorine gas is passed through the
solution for a long time, the precipitate will redissolve. ,

10. Bromine behaves like chlorine except that an excess of this

oxidizing agent does not dissolve the precipitate. (Difference from

lanthanum and didymium.)

11. Cerous salts may be oxidized to ceric salts in acid solutions
(a) by heating with PbOz and HNOg (1:2); (6) by heating with am-
monium persulfate; (c) by electrolysis. In all cases, the solution
_ becomes yellow or orange in color.

2. Ceric Salts

A solution of either ceric nitrate, Ce(NO3)4, or of ceric ammonium
nitrate, Ce(NO3)4-2NH4NO3+H20, should be used.

The beautiful orange-red color of these solutions is characteristic
of all ceric salts, as is also their tendency to form difficultly soluble
basic salts.

THE CERITE METALS 471

Preparation of Ceric Compounds.—As has been already stated, cerous
hydroxide on standing in the air gradually changes to yellow, on account of
the formation of ceric hydroxide. This oxidation takes place immediately
on the addition of chlorine or hypochlorites. If the solution of a cerous salt
is treated with caustic potash solution and chlorine is conducted into it, the
white cerous hydroxide which was at first formed is quickly changed to light-
yellow ceric hydroxide. The latter compound dissolves in dilute acids, forming
orange solutions. It dissolves in concentrated hydrochloric acid with evolution
of chlorine, forming cerous chloride. If white cerous hydroxide is heated in
the air, it loses water and is changed into CeO:, which is nearly white when
cold, dark orange when hot, and is almost entirely insoluble in concentrated
hydrochloric and nitric acids. In the presence of reducing substances (such
~as KI, FeSO,, etc.) it dissolves in acids, forming cerous salts:

2CeO.+8HCl+ 2KI =2KCl+4H,0+1.+2CeCh.

CeO, can also be dissolved by warming with concentrated sulfuric acid,
with evolution of oxygen and formation of cerous sulfate. It can be readily
brought into solution by fusing with potassium pyrosulfate and dissolving
the melt in considerable hot water to which a little acid is added.

If a mixture of cerous and praseodymium hydroxides is ignited in the air,
a cinnamon-colored mass is obtained, which contains all of the cerium as dioxide
and is readily soluble in dilute acids, forming ceric salts. If concentrated
HCl is used, there is an evolution of chlorine, the ceric salt being reduced to
cerous chloride. Concentrated nitric acid dissolves it, forming cerous and ceric
salts; a distinct evolution of oxygen can always be detected.

The reason why the brown mass containing a little praseodymium dissolves
although the pure oxide does not, is probably the following: CeOkz, like MnO,
and PbO, (see pp. 162 and 206), plays the part of an acid anhydride, so that
the brown mass contains the praseodymium as the salt of ceric acid. On
treating this salt with a stronger acid, the praseodymium salt of the latter is
formed, setting free ceric acid (ceric hydroxide), which in the hydrated form
is readily soluble in acids, forming ceric salts.

Basic Ceric Salts.—Ifa solution of ceric nitrate is evaporated on the water-
bath to a consistency of syrup, the mass dissolves readily in water after it has
become cold, and the solution can be boiled without becoming turbid. If,
however, a little nitric acid is added, a yellow precipitate is immediately formed,
consisting of basic ceric nitrate; on the addition of more acid the precipitate
redissolves. This can be explained as follows: By treating the solution of
ceric nitrate with considerable water it becomes hydrolyzed considerably, but
the basic salt produced is present in the hydrosole state and is changed by the
acid into the hydrogele form.

As lanthanum and didymium salts do not yield basic salts under these con-
ditions, this property can be used for separating cerium from these metals.

It is characteristic of cerium to form with ammonium nitrate an easily
crystallizable salt, ceric ammonium nitrate: Ce(NQO;),-2NH,NO;-H,0O.

All ceric salts may be readily reduced by the ordinary reducing ses
(aleohol, HI, SO2, H:S, HNO2, H2O:, ete. ) to cerous salts.

Oxalic Acid added to a concentrated solution of a ceric salt at
first precipitates a dirty orange precipitate which gradually becomes

;

472 REACTIONS OF SOME OF THE RARER METALS

yellow and gelatinous as the addition of oxalic acid is continued, and —
finally crystalline. The precipitate dissolves in a large excess of oxalic
acid, but the solution gradually becomes turbid in the cold, or more —
quickly on heating, as the ceric salt is reduced to cerous salt at the —
expense of the oxalic acid from which carbon dioxide is evolved. Finally
all the cerium will be precipitated as cerous oxalate. (Difference from
La, Di and the yttrium earths.) ;
Ammonium Oxalate behaves similarly.

REACTIONS IN THE DRY WAY

The borax head is colored dark brown when hot and light yellow
to colorless when cold, after being heated in the oxidizing flame. In —
the reducing flame the bead becomes colorless, although strongly
ignited CeOz will remain suspended in the bead, giving it a Bi
yellowish appearance.

LANTHANUM, La. At. Wt. 138.9. M. Pt. 810?

Lanthanum forms only one oxide,* Laz03, which, even after being _
strongly ignited, dissolves readily in acids. Its salts are colorless and

yield no absorption spectrum, so that lanthanum may be distinguished

in this way from didymium and erbium.

REACTIONS IN THE WET WAY

A solution of lanthanum nitrate, La(NO3)3, should be used. +

1. NH:OH and (NH4)2S precipitate a white basic salt which is

difficult to filter. The presence of tartaric acid prevents the precipi-
tation. (Difference from yttrium.)

2. KOH and NaOH precipitate the white hydroxide, La(OH)3.
There is no change to be noticed on treating with oxidizing agents
(difference from Ce). La(OH)3 is soluble enough in water to turn
red litmus-paper blue, and it decomposes ammonium salts on warm-
ing with evolution of ammonia. The fused oxide is realy soluble |
in car

3. (NH4)2CO3 produces a , white vrei slightly soluble in an —
excess of the reagent (difference from aluminium); after standing some
time crystalline lanthanum ammonium carbonate is precipitated.

4. Oxalic Acid produces a white crystalline precipitate, insoluble
an excess of the precipitant and in ammonium oxalate, but soluble in
dilute mineral acids. (Difference from Th.) .

*H.O. is said to cause the formation of LazQ;. Cf. Z. anorg. Chem., 21,
70 (1899).

THE CERITE METALS 473

5. KoSO. precipitates white, crystalline Laz(SO4)3-3K28O4, in-
soluble in a concentrated KeSO, solution.

6. Lanthanum Sulfate is soluble only in ice-cold water; on warm- |
ing the saturated solution to 30° the salt separates out thickly (differ-
ence from cerium.)

7. Iodine —If ammonia is added to a cold, dilute acetic acid
solution of a lanthanum salt, and the slimy precipitate is washed with
water and then treated with solid iodine, the whole mass gradually
assumes a blue color which is similar to that produced by the action
of iodine upon starch (this property is peculiar to lanthanum). The
blue color is destroyed by the addition of acids or alkalies, and dis-
tinguishes it from all the other earths.

8. HF precipitates white gelatinous lanthanum fluoride, which
eventually becomes crystalline LaF3-3H2O; the precipitate is insoluble
in an excess of the precipitant and in dilute acid it is gradually dissolved
by strong mineral acids.

NEODYMIUM, Nd. At. Wt. 143.6. M. Pt. 840°?

yt icgedae PRASEODYMIUM, Pr. At. Wt. 140.5. M. Pt. 940°?

It is very difficult to separate these two metals from one another.
It is accomplished only by repeated fractional crystallization of the
ammonium double nitrates.

Neodymium apparently forms only one oxide, _Nd2Qz; _ it
appears bluish after being ignited, and is readily soluble in acids,
forming violet salts, which afford a characteristic absorption spec-
trum.

Praseodymium, on the other hand, forms a greenish-white oxide,
Pr2O3, which on being ignited is changed into dark-brown peroxide,
Pr4O7. On being heated in a stream of hydrogen, the latter is reduced
back to Pr2O3. The peroxide dissolves in acids with loss of oxygen,
forming greenish salts corresponding to the lower oxide and yielding
a characteristic absorption spectrum.

The Didymium reactions take place with a mixture of the two
elements. A solution of didymium nitrate, Di(NOs)s, is used.

Didymium salts are violet and show a characteristic absorp-.
tion spectrum (difference from Ce and La). The behavior toward
NH4OH, (NH4)2S, KOH, (NH4)2COsz, and KeSOs is exactly the same
as with lanthanum. Oxalic acid precipitates the reddish sige
which in other respects is like lanthanum oxalate.

Absorption Spectrum of Praseodymium.—596.9up, 5904yu, 482.1upy,
469.5uyu, and 444.1yu. (Continued on p. 480.)

REACTIONS OF SOME OF THE RARER METALS

474

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SHLYUVA AHL JO SNOILOVaEY AHL JO AYVWWOAS

475

THE CERITE METALS

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REACTIONS OF SOME OF THE RARER METALS

476

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(‘panuyuo))—SHLUVA GAUVU AHL AO SNOILOVaY HHL JO AUVWNOAS

477

THE CERITE METALS

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REACTIONS OF SOME OF THE RARER METALS

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478

479

THE CERITE METALS

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480 REACTIONS OF SOME OF THE RARER METALS
Absorption Spectrum of Neodymium.—T29.1yp, 690.6up, (579.7up,

675.9uu), 531.7up, (522.2uu, 520.9uu), 512.0uu, 509.6uy, 482.1yy,
475.9up, 469.5uu, 461 4up, 444.3, 434.1lup, 427.7uy, 417.3up.

TANTALUM, Ta. At. Wt. 181, and NIOBIUM, Nb. At. Wt. 93.5

These two rare elements, belonging to the nitrogen-vanadium group,
form oxides of the formula R2Os5, which behave as acid anhydrides

and should have been considered, perhaps, under the acids. Since, —

however, tantalic and niobic acids are soluble, under certain conditions,

in strong acids and from these solutions are precipitated by ammonia —

and ammonium sulphide, it seems better to consider them at this place.

Occurrence.—In the form of meta-acids these elements appear in
the isomorphous minerals tantalite, Fe(TaOg3)2 and niobite or columbite,
Fe(NbO3)2. In tantalite a part of the tantalic acid is replaced by
niobic acid and a part of the iron by manganese. Niobite shows an
analogous behavior.

In the form of pyro acids the two elements occur as an isomorphous
mixture in the mineral yttrotantalite, Y4(Ta2O7)3 and Y4(Nb207)3.
Finally tantalum, and to some extent niobium, replaces the phos-
phorous’ in monazite; (Ce,La,Di)POs. Tin is usually found in all

the above minerals and often tungsten; conversely cassiterite and —

wolframite often contain small quantities of niobic and tantalic acids.

TANTALUM, Ta. At. Wt. 181. M. Pt.=2850°. Sp. Gr.=16.5

Metallic tantalum* is ductile, although the presence of a little
impurity makes it harder than tool steel. On ignition in the air it —

assumes a yellow to blue tinge caused by a thin coating of oxide.
Tantalum is not attacked by boiling H2SO4, HCl, HNOs, or even aqua

regia, but it is slowly dissolved by hydrofluoric acid with evolution —
of hydrogen; any metal remaining undissolved is then brittle on —
account of absorbed hydrogen. The concentrated solution in hydro- —
fluoric acid forms with concentrated KOH insoluble, crystalline potas- —
sium tantalum fluoride, KoTaF7. By evaporating the solution in —
HF with concentrated H2SO,4 until the former acid is all expelled, the —
residue dissolves in a little cold water, but the solution becnaas turbid |

on dilution or especially by boiling.
Tantalum forms two oxides, Ta2O4 and Ta2QOs, the former being

indifferent chemically and the latter an acid anhydride. After ignition, —
the pentoxide is insoluble in acid and is not rendered soluble by fusion —

*W. von Botton, Z. Electrochemie, 11, 45 (1905).

7 << wre

THE CERITE METALS 481

with pyrosulfate, although it is volatilized by heating with ammonium
fluoride. Fusion of the oxide with caustic alkali in a silver crucible
gives rise to alkali tantalates, both the meta- and hexatantalates being
known; only the former are soluble in water.

Potassium hexatantalate, KeTagO19+16H20, is soluble in water
and caustic potash solution, while the sodium salt is soluble in water,
but not in caustic soda. If potassium hexatantalate is treated with
hot water, a part of it goes into solution and undergoes hydrolysis,
forming a colloidal solution of HgsTagO19. If COe is conducted into the
solution, the tantalic acid is completely precipitated in a flocculent
condition. (Difference from niobic acid.) The remaining tantalates
are all insoluble.

REACTIONS IN THE WET WAY

A solution of potassium hexatantalate should be used.

1. Mineral Acids.—(a) H2SO.4 precipitates tantalic acid from
cold, dilute solutions, and the precipitation becomes nearly quantita-
tive on boiling. Hot, concentrated H2SOx4 dissolves the precipitate
produced by the dilute acid. On diluting the solution with water
after it has become cold, the tantalic acid is reprecipitated (difference
from niobium.)

(b) HCl added to a concentrated solution at first produces a pre-
cipitate, which dissolves in an excess of the acid, forming an opalescent
solution. From this solution sulfuric acid precipitates tantalic acid
in the cold, but the precipitation is not quantitative even on boiling.

2. NH4OH and (NHsgz)2S precipitate from the hydrochloric acid
solution either tantalic acid itself or an acid ammonium tantalate;
tartaric acid prevents the precipitation.

3. NaeCOz produces a .partial precipitation of tantalic acid when
added to an acid solution of a tantalate, but the precipitate dissolves
in an excess of the precipitant.

4. HeS precipitates tantalic acid in the cold, especially from sul-
furic acid solution, and the precipitate is almost quantitative.

5. H2O2 dissolves freshly precipitated tantalic acid if acid or
alkali is present. From the solution thus obtained, the tantalic acid
is not precipitated by the above reagents unless the hydrogen peroxide
is destroyed by boiling the alkaline solution or by the action of sulfurous
acid.

6. Ks[Fe(CN)6| and KCNS produce white precipitates.

7. Tincture of Nutgalls produces no precipitate (difference from
niobic acid). |

482 REACTIONS OF SOME OF THE RARER METALS

8. K4[Fe(CN).] produces in acid solutions a light-yellow precipi- —
.. tate, which becomes brown on the addition of a little ammonia.

9. H{KF>2|.—If a concentrated solution of tantalic acid in hydro-
fluoric acid is treated with KF, the difficultly soluble Keo{TaFz]. is
formed, which separates from the solution in the form of orthorhombic

needles (200 parts of water dissolve 1 part of salt) difference from

niobium). On boiling the solution of tantalic potassium fluoride, the
very difficultly soluble oxyfluoride precipitates (K4TasOsFi4). By
means of this reaction the merest trace of tantalic acid can be detected
in the presence of niobic acid.

10. Zn and HCl do not produce colored solutions (difference from
niobium).

REACTIONS IN THE DRY WAY

Ta2Os5 is infusible. The bead of salt of phosphorus remains

colorless in both oxidizing and reducing flames. The addition of

FeSO. does not cause the formation of a blood-red color. (Difference
from Ti and Nb.)

NIoBIUM. Sp. Gr.=12.7. M. Pt.=1950° C.*

The metal niobium is very similar to tantalum; it is more readily

attacked by acids.

Niobium forms three oxides: Nb2O2, Nb204, and Nb2Os, of which
the last is an acid anhydride. Nb2Qs, like Ta2Qs, is insoluble in acids
after it has been ignited and is not rendered soluble by fusing with
potassium pyrosulfate. The melt dissolves in cold water, but niobic
acid separates from the solution on boiling. By fusing with KOH or
KeCO3, potassium hexaniobate, KgNbeg0i9:16H20, is formed, which
is soluble in water. The corresponding sodium salt is insoluble in
caustic soda solution, but soluble in water.

Sodium hexaniobate is largely hydrolyzed in aqueous solution, form-
ing a colloidal solution of niobic acid, HgNbeOi9. By passing CO2
through this solution for half an hour, there is no precipitation of
niobic acid, but only after long standing. (Difference from tantalic
acid.)

REACTIONS IN THE WET WAY

Use a solution of potassium hexaniobate.

1. Mineral Acids produce in alkali niobate solutions a white,
amorphous precipitate of niobic acid, which is only slightly soluble
in an excess of the acid. Concentrated sulfuric acid, however, dissolves

* WERNER VON Bouton, Chem. Zentralbl., 1905, I, 586.

re ae

THE CERITE METALS 483

the niobie acid on warming; and the solution remains clear after being
diluted with cold water. (Difference from tantalum.)

By boiling the diluted solution, the niobic acid is almost com-
pletely precipitated, but in a very finely divided condition, such that
it is hard to filter. If the acid is exactly neutralized with ammonia,
the niobic acid is precipitated in a flocculent condition easy to filter.
It is best to wash such a precipitate with 0.5 per cent ammonia water
or with 1 per cent acetic acid. Washing with pure water causes a
turbid filtrate and mineral acids should not be used.

If a solution of niobic acid in sulfuric acid is poured into a con-
centrated solution of ammonium sulfate, no niobic acid is precipitated
by boiling. (Difference from tantalic acid.)

If the niobic acid is treated with boiling hydrochloric acid, it dis-
solves only slightly, but on pouring off the acid, the residue is soluble
in water.*

Carbonic acid decomposes sodium niobate to some extent. On
the other hand, a niobate solution prepared after fusing with sodium
carbonate and potassium nitrate is not decomposed by COe. (Differ-
ence from tantalic acid.)

2. NHiOH and (NH4)2S precipitate niobic acid from the sul-
phuric acid solution, and the precipitate is soluble in HF.

3. H2O2 reacts as with a tantalate.

4. Tincture of Nutgalls produces no precipitate.

5. K4{Fe(CN).] produces a grayish-green precipitate.

6. H{[KF2].—If niobic acid is dissolved in an excess of HF and
KF is then added, readily soluble niobic potassium fluoride is formed
(12.5 parts of water dissolve 1 part of the salt). By boiling the dilute
aqueous solution, soluble potassium niobic oxyfluoride is formed, which
is even more soluble. (Difference from tantalum.)

7. Zine produces in an acid solution of a niobate a dirty-blue color-
ation which disappears after some time. (Difference from tantalum.)

REACTIONS IN THE DRY WAY

The bead of salt of phosphorus is blue, violet, or brown in the
reducing flame (according to the amount of niobic acid which is
present); the bead becomes red on the addition of FeSOs.

* This behavior reminds one of metastannic acid. (Cf. p. 257.)

484 REACTIONS OF SOME OF THE RARER METALS

Separation of Tantalum from Niobium a

* ey
ae,

According to Weiss-Landecker* and Hauser-Lewite.t | =a

Fuse the two oxides with a little sodium carbonate in a platinum crucible
cool, extract with hot water and filter off the undissolved sodium hexatanalate. —
Wash the residue with a solution of NaHCO; and conduct CO; into the filtrate, —
whereby some flocculent tantalic acid will be precipitated, but the niobic acid —
will remain in solution. Filter, unite this precipitate with the residue from —
the sodium carbonate fusion and dissolve both in sulfuric acid and hydrogen — 4
peroxide. Pass SO, into this solution and boil; tantalic acid will be precipi- —
tated. Similarly saturate the sodium carbonate solution of sodium niobate —
with SO, and niobic acid will be precipitated by boiling. aa

- ey
Vig

*Z. anorg. Chem., 64, 65-103 (1909).
7 Z. angew. Chem., 1912, 100.

METALS OF THE H2S GROUP

THALLIUM, VANADIUM, MOLYBDENUM, TUNGSTEN, SELENIUM,
TELLURIUM, RHODIUM, PALLADIUM, OSMIUM, IRIDIUM, RU-.
THENIUM

THALLIUM, Tl. At. Wt. 204.0. Sp. Gr.=11.9. M. Pt.=302°

Occurrence.—Thallium is found in nature very sparingly; in small
amount in many varieties of pyrite, and accompanying potassium in
carnalite and sylvite, in many lithium micas and in many mineral waters.
It replaces the silver to a considerable extent in copper-silver selenide,
in crookesite, (AgT]Cu)2Se, and in berzelianite, (CuAgTl)2Se. There
are no characteristic thallium minerals. The principal sources of our
thallium is the dust from sulfuric acid plants where pyrite containing
thallium is used.

Metallic thallium reminds one of lead in its color, softness, high
specific gravity, and low melting-point.

Thallium dissolves readily in nitric and sulfuric acids, but not in
hydrochloric acid. It forms two oxides: thallous oxide, Tl2O, and
thallic oxide, Tl,03; both are anhydrides of bases and from them
thallous and thallic salts are derived.

REACTIONS IN THE WET WAY

A. Thallous Compounds

Thallous compounds are colorless and soluble in water as a rule. The sul-
fide, chloride, bromide, iodide, and chromate are insoluble in water. Thallous
oxide is a colorless powder, whose aqueous solution reacts alkaline and absorbs
carbon dioxide with avidity.

Use a solution of thallous sulfate for the following reactions:

1. HeS causes no precipitation from solutions which contain
mineral acids; in neutral solutions, thallium is incompletely precipi-
tated as black thallous sulfide, TleS. ToS is readily soluble in mineral
acids, but insoluble in acetic acid and alkaline sulfides. It is oxidized
readily on standing in the air to thallous sulfate.

2. (NH4)2S precipitates all of the thallium as TlS.

3. KOH, NaOH, or NH4OH produces no precipitation.

485

486 REACTIONS OF SOME OF THE RARER METALS

4, Alkali Carbonates cause precipitation only in very concen-
trated solutions, for thallous carbonate is fairly soluble (100 parts of
water dissolve 5 parts of the salt.)

5. HCl produces a heavy, white precipitate of thallous chloride, —

very slightly soluble in water, and still less so in water containing a
little hydrochloric acid. 7
6. KI precipitates yellow thallous iodide, Tl, from even the most —
dilute solutions; this is the most sensitive reaction for thallium.
7. Alkali Chromates precipitate yellow thallous chromate, insoluble
in cold nitric or sulfuric acids.

8. Ho[PtCle] precipitates light-yellow thallium chloroplatinate,
which is almost entirely insoluble in water; 1 part dissolves in 15,600
parts of water at 15° C. and in 1950 parts of water at 100° C.

9. Alo(SO4)3.—If a solution of thallous sulfate is treated with

aluminium sulfate and the solution is then allowed to crystal- 4
lize, glistening, colorless octahedrons are obtained of thallium alum,

TIAI(SO4)2+12H20. |
10. K3[Fe(CN)6] precipitates brown T1(OH)s3 in alkaline solutions:

2[Fe(CN)¢|=+30H~+TI* — 2[Fe(CN)6]=+T1(OH)s.

Thallium is like lead in respect tc its specific gravity and to the solubility

of its halogen compounds; but, on the other hand, it is similar to the alkalies
with regard to the solubility and alkaline reaction of the hydroxide and car-—
bonate, and with regard to its forming an insoluble chloroplatinate and an
alum.

B. Thallic Compounds

Thallic compounds cannot as a rule be prepared by the oxidation of thallous
compounds (with the exception of thallic chloride, which is readily obtained
by the action of chlorine water upon thallous chloride). They are obtained
by the solution of thallic oxide * in acids, and can be distinguished from thallous
compounds by the ease with which they suffer decomposition in aqueous
solution. Thus thallic sulfate is decomposed on boiling its aqueous solution
into thallic hydroxide and sulfuric acid; the nitrate behaves similarly.

The chloride, TICl;, is a hygroscopic and not very stable substance; on
being heated to 100° C. chlorine is evolved with the formation of thallous
chloride.

1. KOH, NaOH, and NH,OH8 precipitate brown thallic hydroxide,
TI(OH)3, from solutions of thallic salts, which changes to TIO(OH)
on standing in the air, is difficultly soluble in acids and insoluble in
an excess of alkali.

* T1.O; is not attacked in the cold by concentrated sulfuric acid, but is dissolved
on warming. The hydrated oxide, TIO(OH), is much more soluble.

METALS OF THE H.S GROUP 487

2. HCl and alkali chromates do not cause precipitation.
3. KI precipitates thallous iodide with deposition of iodine.

REACTIONS IN THE DRY WAY

Thallium salts color the non-luminous gas-flame a beautiful emerald
green. The thallium spectrum consists of a green line at 535.0, nearly
coincident with the green barium line at 534.7.

VANADIUM, V. At. Wt. 51.0. M. Pt.=1730°

Occurrence.—Vanadinite, Pbs(VO4)3Cl; carnotite; * mottramite,
(CuPb)2V2010+2H20; many clays and in almost all granites.

| Vanadium, like nitrogen, forms five oxides: V20, V2O2, V20s, V2O0a,

V20s.

The first three of these oxides are basic anhydrides. Compounds
representing these valencies of vanadium are not encountered in quali-
tative analysis except, to some extent, in the tests for vanadium with
strong reducing agents. The oxides V2O4 and V20Os represent the types
of vanadium compounds usually encountered in analytical chemistry.

V204 is the anhydride of hypovanadic acid, V202(OH)4. This com-
pound is an amphoteric substance and forms salts with both acids and
bases. V2O« itself is a blue powder, soluble in concentrated acids,
forming blue divanady] salts:

V204+2H2S04 = V202(S04)2+2H20.

If the solution of divanadyl] sulfate is treated with sodium carbonate
or ammonia (avoiding excess), hypovanadic acid separates out as a
grayish-white precipitate, which, like the anhydride, is soluble in acids
with blue color and in alkalies with a brown color. The alkali hypo-
vanadates correspond to the symbols NazV205 and NazV4O9. Hypo-
vanadates of other metals are for the most part insoluble in water and
such precipitates may form when an acid solution containing vanadium
is neutralized. The divanadyl compounds are readily formed by
reducing solutions of the pentoxide in mineral acids with sulfurous acid
(cf. p. 489), and serve, on account of their blue color, for the detec-
tion of vanadium.

V20s is the anhydride of vanadic acid and is an orange-red crystal-
line mass, which is readily fusible but non-volatile. It is only slightly

* According to FrrepeL and CuMENGE, carnotite contains 18 per cent V2O; and
55 per cent VOs, as well as K, Ca, Ba, H, As, and P. (Cf. HiLLEBRAND and Ran-
soME, Am. J. Science, 10, 138.)

488 REACTIONS OF SOME OF THE RARER METALS

soluble in water; forming a slightly acid, yellow solution, but readily —
soluble in concentrated solutions of caustic alkalies, forming vanadates. —
Like phosphoric acid, vanadic acid exists in the form of meta-,
pyro-, ortho-, and poly-compounds, of which the meta-compounds are ~
the most stable and the ortho-compounds the least so. Thus an aqueous 7
solution of potassium or sodium orthovanadate is hydrolyzed, even in —
the cold, into the pyro-salt and alkali hydroxide,

2Na3VOs+H20 @ NasV207+2Na0H,
and on boiling the meta-salt is formed:
Na4V207+ H20 @ 2NaVO3+2Na0H.

The meta-, pyro-, and ortho-salts of the alkalies are colorless or
slightly yellow, while the polyvanadates, e.g., the tetra- and hexa-— a
vanadates, are intensely orange or reddish. Thus the colorless or —
light-yellow solutions of the ortho-, meta-, and pyrovanadates are —
colored intensely orange on the addition of acid.

Besides the above types of vanadium compounds, this element 4
exists as pervanadic acid, HVO,, formed by the addition of hydrogen — 4
peroxide to the acid solution of a vanadate. 3

The reactions of quadrivalent and quinquevalent vanadium will be
considered together; the other forms are not common enough to make” 7
it necessary to describe their characteristic reactions. .

REACTIONS IN THE WET WAY r

1. NH,Cl.—If a piece of solid ammonium chloride is added to a —
solution of an alkali vanadate, colorless ammonium metavanadate —
separates out,

NasV207+4N H4Cl=2NH4V03+2NH3+H20+4NaCl,

difficultly soluble in a concentrated solution of ammonium chloride.

2. Pb(C2H302)2 precipitates vanadic acid quantitatively as vheme zs
lead vanadate: .
3Pbt *+2VO04 > Pb3(VO4)e.

This precipitate, however, is more soluble in dilute nitric acid than is lead i
chromate, and it is possible to separate chromic acid from vanadiec acid by —
treating the solution of the two acids with lead nitrate in dilute nitric acid 4
solution; under the proper conditions all but a fraction of a milligram of the
chromium is precipitated and 100 mgms. of vanadium yield no precipitate.* —
The vanadium can be detected in the filtrate by the HO: test. a

* The solution is neutralized exactly with NaOH, 2 cc. of HNO; (sp.gr. 1 20)
are added together with an excess of 20 per cent Pb(NOs)2 solution.

METALS OF THE HS GROUP 489

3. NH,OH added to a solution of an alkali vanadate causes no
precipitation. Vanadic acid behaves like phosphoric acid toward
ammonia; vanadates of ferric iron, aluminium and uranium are likely
to be precipitated by ammonia, as well as vanadates of the alkaline
earth metals.

NH4OH added to a solution of a vanadyl salt precipitates dark-
gray hypovanadic acid:

V202Cl4,+4NH40H =4N H4Cl4+ V202(OH)a.

The precipitation is not quantitative and small quantities of vanadium
may remain in solution when the vanadyl salt is alone present. If, however,
an excess of ferric chloride is added to the solution, the vanadium is quan-
titatively precipitated upon the addition of ammonia. This is true both of
vanadie acid and of vanadyl salt, ferric vanadate and ferric hypovanadate
being precipitated with the ferric hydroxide. Instead of ferric chloride, an
aluminium or uranium salt may be used for the same purpose.

4. (NH4)2S produces no precipitation, but causes the solution
to turn brown, owing to the formation of thio-salts.

If hydrogen sulfide is conducted into a strongly ammoniacal solution of
a vanadate or hypovanadate, the solution at first turns yellowish-red, but the -
color slowly deepens until eventually a characteristic, brilliant, violet-red color
is obtained when the solution has become saturated with H.S. Ammonium
salts interfere somewhat with this test, but their influence is overcome by the
addition of a large excess of ammonia. The red color is probably caused by
the formation of ammonium thiovanadate. As little as 0.2 mgm. of vanadium
can be recognized by the red color.

The addition of acid to the red solution produces a black precipitate of
V8, or V.S;. The precipitation is not quantitative; the filtrate is always
colored blue and contains detectable amounts of vanadyl salts. The precipi-
tate is soluble in alkalies, alkali carbonates, and in alkali sulfides, forming a
brown solution.

Molybdenum gives a similar red color, in case it was not completely removed
by previous treatment with hydrogen sulfide in acid solution, and obscures the
above test or may be mistaken for vanadium.

5. H2S gives no precipitation in acid solution, but reduces com-
pounds of vanadic acid to divanadyl compounds, so that the solution is
colored blue:

2H3VO4+H28+4HCl= V202Cl4+6H20-45.

6. Reducing Agents (SO2, H2S, HBr, alcohol, oxalic and tartaric
acids, sugar, etc.) reduce acid solutions containing vanadates to blue
vanadyl salts:

pee ee
2V04=+S037 +10H*t —> [V202]* +4S$0,-+5H.20.

490 REACTIONS OF SOME OF THE RARER METALS

HI reduces vanadic acid to green salt of V20s3:
V0O,*+21-+-8H* — Vit t4To +4H.20.

The green color only appears after the iodine has been removed by —
continued boiling of the solution.

Metals, such as Zn, Al, and Cd, cause still further reduction of vanadie
acid, so that the solution turns at first blue, then green, and finally violet. j

Boiling an acid solution of a vanadate with concentrated hydrochloric acid
and alcohol reduces the vanadium quantitatively to divanadyl salt. Treat-
ment of vanadic acid with ferrous salt also reduces the former to divanadyl —
salt and the excess of the ferrous iron can be oxidized by cold potassium di- — 4
chromate solution without oxidizing the vanadium. a

7. H2O2.—If an acid solution of a vanadate is treated aie a
few drops of H2O2 and shaken, the solution becomes colored reddish
brown owing to the formation of pervanadic acid, HVOs, insoluble 1 In F
ether. This is a very delicate reaction.

8. Mercurous Nitrate precipitates white mercurous vanadate from 4
neutral solutions of a vanadate; the precipitate is soluble in nitric acid. —

9. Oxidizing Agents convert divanadyl compounds into vanadic —
acid. The oxidation may be effected by bromine in hydrochloric — ;
acid solution,

[V20o]**-+Bro-+6H20 — 2V0."-+2Br-+ 12H", .

by dilute potassium permanganate in hot, very dilute sulfuric acid : .
solution,

++ |
5[V202]* t+2Mn0.-+22H20 > 10VO4= +2Mn*+++44Ht,

or by sodium peroxide in alkaline solution:

++
[V202]* ++ Na202+8 OH- > 2VO04=+2Nat+4H20.

Chromium, aluminium, vanadium and uranium may be separated from iron,
nickel, cobalt and manganese by means of this reaction; the chromate is left —
in solution as sodium chromate, NazCrO,, the aluminium as sodium aluminate,
NaAlO,, the vanadium as sodium vanadate, Na;VO,, the uranium as sodium —
peruranate, and the zinc as sodium zincate, Na2:ZnO,; while the iron is pre- —
cipitated as Fe(OH);, the nickel as Ni(OH). or Ni(OH);, the cobalt as Co(OH)s, —
and the manganese as hydrated MnO,. <j

Detection of Vanadium in Rocks (Hillebrand) * q

Fuse 5 gms. of the finely powdered rock with 20 gms. of Na,CO; and 3 gms. —
of NaNO;. Cool, extract the fused mass with water, reduce the manganate —
formed by the addition of a little alcohol, and filter the solution which contains _

* Amer. J. Science, 1898, p. 209.

METALS OF THE HS GROUP 491

the sodium salts of arsenic, phosphoric, molybdic, chromic, vanadic and tungstic
acids. Nearly neutralize it with nitric acid (the amount necessary having been
determined by a blank test), evaporate nearly to dryness, take up in water,
and filter. Treat the alkaline solution with mercurous nitrate, whereby mer-
curous phosphate, arsenate, chromate, molybdate, and tungstate with some
basic mercurous carbonate are precipitated. Boil the solution, filter, dry the
precipitate, separate it from the filter, ignite it in a platinum crucible, and
fuse with a little sodium carbonate. Extract the fused mass with water, when
a yellow color shows that chromium is present. Acidify the solution with
sulfuric acid, and precipitate traces of Pt, Mo, and As by means of HS (best
in a small suction flask). Filter off this precipitate and remove the excess of
H.S from the filtrate by boiling, while passing a stream of carbonic acid gas
through it. Evaporate the solution to dryness, and carefully expel the excess
of sulfuric acid, heating in an air-bath. Dissolve the residue in 2 or 3 cc. of
water and shake with a few drops of H.O.; a brownish-yellow color shows
the presence of vanadium. If chromium is present, on adding H.O, and ether
to the sulfuric acid solution and shaking, the ether will be colored blue by
chromium and the aqueous solution yellow by vanadium.*

REACTIONS IN THE DRY WAY

The borax bead is colorless in the oxidizing flame if slightly saturated
with the vanadium compound, yellow if strongly saturated, and green
in the reducing flame.

MOLYBDENUM, Mo. At. Wt. 96.0. M. Pt.=2500?

Occurrence—Molybdenite, MoSe; wulfenite, PbMoQ.; powelite,
CaMoO4. Molybdenum has a valence of 2, 3, 4, and 6, and forms the
following oxides: MoO, Mo203, MoOe2, and MoOs. The first three
are basic anhydrides, while the last oxide, MoOs, is an acid anhydride,
forming a white mass (yellow when warm) which is readily fusible,
but very difficultly volatile. When heated strongly, colorless, trans-
parent, thin, orthorhombic plates of MoO3 may be obtained from
the fumes. MoQOgz is only very slightly soluble in water, but dissolves
readily in alkalies and in ammonia, forming molybdates. Molybdic
acid itself can be readily obtained as a solid mass by acidifying the
solution of an alkali molybdate; it is soluble in an excess of the acid
(difference from tungstic acid). . The most important commercial
molybdate is the acid ammonium molybdate, corresponding to the
formula:

(NH4) 6Mo7O24 +4H20.

*E. Cuampaane, Chem. Zentralbl., 1904, II, p.-1167.

492 REACTIONS OF SOME OF THE RARER METALS

REACTIONS IN THE WET WAY

A solution of ammonium molybdate should be used.
The alkali molybdates are soluble in water; the remaining salts are mostly '
insoluble in water but soluble in acids.

1. Dilute Acids precipitate from concentrated alkali molybdate —
solutions white HzMoO,, soluble in an excess of acid. j
Concentrated Sulfuric Acid.—If a trace of a molybdenum com- ~
pound is evaporated with a drop of concentrated sulfuric acid almost —
to dryness in a porcelain dish, the mass is colored intensely blue. This —
is an exceedingly delicate reaction.
2. H2S at first colors acid molybdenum solutions blue, and pre- 7
cipitates, little by little, the molybdenum as brown molybdenum tri- —
sulfide, MoSs, soluble in ammonium sulfide, forming a brown solution —
from which MoS83 ‘is reprecipitated by the addition of acids. Molyb- ~
denum sulfide is oxidized by treatment with concentrated nitric acid, —
or by roasting in the air, into MoOs. ¥
3. Zinc.—If a molybdate solution which is acid with hydrochloric —
or sulfuric acid is treated with zinc, the solution is colored at first blue, —
then green, and finally brown. Other reducing agents such as SnCle,
Hge(NOs3)2, etc., cause the same reaction.
4. SOs does not reduce dilute, strongly acid solutions of molyb- —
dates either in the cold or on heating. Neutral or slightly acid
solutions are reduced and colored blue. “@
5. KCNS causes no change when added to acid molybdenum —
solutions, but if the solution is then treated with zine or stannous
chloride, a blood-red coloration is produced on account of the forma-
tion of molybdenum thiocyanate; the reaction also takes place in the —
presence of phosphoric acid (difference from iron). If the solution —
is shaken with ether, the colored compound is dissolved in the latter. ©
6. Sodium Phosphate.—If a few drops of a solution of sodium ~
phosphate are added to a molybdate solution strongly acid with nitric
acid, a yellow crystalline precipitate of ammonium phosphomolybdate \
is formed, slowly in the cold, but much more quickly on warming the
solution (cf. Phosphoric Acid, p. 379). Arsenic acid causes the pre- 4
cipitation of a similar compound (ef. p. 231).
7. Mercurous Nitrate precipitates white mercurous molybdate E
from neutral solutions; the precipitate is soluble in nitric acid. 2
8. Lead Acetate pile, eg white lead molybdate, soluble in 4
nitric acid. a
9. Potassium viecsteahte produces a reddish brown precipitate 4
A very sensitive test. 4

A a

METALS OF THE H,S GROUP 493

Molybdenum solutions containing free oxalic, acetic or phosphoric acids
usually give no precipitate with potassium ferrocyanide, but merely a brown
coloration. Molybdenum ferrocyanide is, however, insoluble in dilute mineral
acids but is dissolved by concentrated hydrochloric acid and reprecipitated
upon diluting. It is readily soluble in caustic alkali and ammonia solutions,
in which respect it is different from the uranyl and cupric ferrocyanides (pp. 156
and 220). To detect the ferrocyanide ion in molybdenum ferrocyanide,
dissolve the salt in ammonia, saturate the ammoniacal solution with H,S,
acidify with dilute H.SO,, filter off the MoS; and test the filtrate with ferric
chloride solution (p. 150).

10. H.O;. If a solution to be tested for molybdenum is evaporated to
dryness on the water-bath, the residue treated with a little concentrated
ammonia and then with hydrogen peroxide, the ammoniacal solution is immedi-
ately turned- pink or red. Then, evaporating to dryness again and treating
the residue with sulfuric or nitric acid, yellow permolybdic acid, HMoOQ,,
is obtained.

REACTIONS IN THE DRY WAY

Alkali molybdates, alone or with sodium carbonate, are reduced
on charcoal to gray molybdenum, a white incrustation of MoOg3 being
formed at the same time.

Salt of Phosphorus Bead.—All molybdenum compounds color the
bead, but the color depends upon the concentration. In the oxidizing
flame the hot bead is colored brownish-yellow to yellow; it becomes
yellowish-green on cooling and finally colorless... In the reducing flame
the bead becomes dark brown when hot and grass green when cold.
The borax bead is similar but not quite as characteristic.

Separation of Molybdenum, Arsenic, Antimony, and Tin

These elements are all precipitated as sulfides upon the introduction of
hydrogen sulfide into an acid solution. They are separated from the members
of the copper group by treatment with ammonium polysulfide solution, in which
their sulfides are soluble. On acidifying this solution of the thiosalts with
dilute hydrochloric acid, the molybdenum, arsenic, antimony and tin are repre-
cipitated as sulfides.

Filter off this precipitate, wash, dry and introduce it, little by little, into a
nickel crucible containing a molten mixture of 10 parts NaOs, and 10 parts
Na,CO; for each part of sulfide precipitate. Fuse the contents of the crucible,
after all the sulfide has been introduced, for ten minutes over the Bunsen
burner, then cool and extract with cold water. The aqueous solution thus
obtained may contain sodium arseniate and sodium molybdate and the insoluble
residue may consist of sodium antimonate and tin dioxide. Filter off this
residue and wash it with a normal solution of sodium hydroxide. Test the
filtrate for arsenic by acidifying with hydrochloric acid, making strongly
ammoniacal and adding magnesium mixture. A white crystalline precipitate
of magnesium ammonium arseniate is formed if arsenic is present, but only

494 REACTIONS OF SOME OF THE RARER METALS

after standing for some time with a little arsenic. Saturate the filtrate from

the magnesium ammonium arseniate, precipitate with hydrogen sulfide to con-
vert any molybdenum present into ammonium, thiomolybdate, and then —

acidify the solution with dilute hydrochloric acid. Filter off the precipitated
sulfide, treat it with concentrated nitric acid in a porcelain crucible and test
for molybdenum with concentrated sulfuric acid, as described on p. 492.

Test for Antimony and Tin.—Treat the residue insoluble in dilute caustic
soda solution with a mixture of equal parts concentrated hydrochloric acid and
water, place the solution in contact with a piece of platinum foil and test for
antimony and tin as described on p. 275.

TUNGSTEN, W. At. Wt. 184. M. Pt.=3000°

Occurrence.—Tungsten is not very often found in nature, but there
are a number of well-crystallizing tungsten minerals, such as the
minerals of the Scheelite group.

Scheelite, CaWO.; cuproscheelite, (CaCu)WO4; reinite, FeWO,;
stolzite, PoWO.4. These minerals all crystallize in the tetragonal system
and form with powellite, CaMoO., and wulfenite, PbMoO., a very
interesting isomorphous group. Another isomorphous group, which

consists of minerals crystallizing in the monoclinic system, is formed —

by hiibnerite, MnW0O.,; wolframite, (MnFe)WO.a, and _ferberite,

FeWO,4. The most important tungsten mineral is wolframite, which

is usually contaminated with small amounts of silicic, tantalic, and
niobic acids. Tungsten forms two oxides, WO2 and WOQOs.

WOz is a brown powder, readily obtained by heating WOz3 to dull
redness in a stream of hydrogen. It is pyrophoric and must, therefore,
be cooled in a stream of hydrogen before it is allowed to come into
contact with the air. By igniting strongly in a stream of hydrogen,
metallic tungsten is obtained, which is stable in the air. This behavior
is important and is taken advantage of in the quantitative deter
tion of tungsten.

WOs is an acid anhydride obtained by the ignition of tungstic acid
of ammonium or mercurous tungstates, or by the oxidation of the
dioxide on heating in the air.

The trioxide is a canary-yellow powder, insoluble in water and
dilute acids, and only slightly soluble in concentrated hydrochloric and

hydrofluoric acids. It dissolves readily by warming with potassium

or sodium hydroxides, and less readily in ammonia. It is most easily
dissolved by fusing with sodium carbonate, sodium tungstate being
formed:

W03+ Na2CO3 = NazW04+COrz.

ee ee ee ee eee

ev 2 a

METALS OF THE H.S GROUP * 495

It is changed to potassium tungstate by fusing with potassium
pyrosulfate:
W0O3+ K28207 = K2W04+ 2803.

If the product of this last fusion is treated with water, usually none
of the tungsten goes into solution, because if an excess of potassium
pyrsoulfate is present (which is usually the case) it reacts with the
potassium tungstate, forming free tungstic acid:

K2eW04+ Ke8207+ H20 = 2K2804+HeWOg.

If not enough pyrosulfate remains to complete the above decom-
position, some of the tungsten will be dissolved, but never all of it.
If a little sulfuric acid is added to the water, none of the tungsten
will go into solution. This property enables one to separate tungsten
from titanium. If ammonium carbonate is added, all of the tungsten
dissolves, which enables us to separate tungstic from silicic acid.

REACTIONS IN THE WET WAY

A solution of sodium tungstate should be used.

1. Mineral Acids, HCl, HNOz, H2SO.4, produce, in the cold, a
white, amorphous precipitate of hydrated tungstic acid, H2W0O4+ H20.*
By boiling the solution, the yellow anhydrous acid HzWOs is obtained,
insoluble in dilute acids, but soluble to an appreciable extent in con-
centrated hydrochloric acid.

Tungstic acid must always be washed with water which contains
acid or a dissolved salt, as otherwise tungstic acid will form a pseudo-
solution with pure water, so that a turbid filtrate will be obtained
(ef. pp. 58, 127 and 218).

Phosphoric acid behaves differently toward solutions of the alkali
tungstates than do the other mineral acids; it produces a white pre-
cipitate soluble in an excess of phosphoric acid; a complex phospho-
tungstic acid is formed, e.g., NagPO4-12WO3. If the solution of an
alkali tungstate is boiled with free tungstic acid, the latter gradually
goes into solution, forming a metatungstate;

NaeWO,4 +3WOz3 = NaeW.40i13.

Mineral acids cause no precipitation in solutions of metatung-
states. If the solution is boiled with an excess of acid, the soluble
metatungstic acid is gradually changed to insoluble, ordinary tungstic
acid, which is then precipitated.

* The presence of tartaric acid prevents the precipitation.

496 REACTIONS OF SOME OF THE RARER METALS

2. HeS produces no precipitation in acid solutions. .
3. (NH4)2S gives no precipitation in a solution of an alkali tung- |
state, but if the solution is afterward acidified, light-brown tungsten —
trisulfide, WS3, is precipitated, which has the property of forming —
pseudo-solutions with pure water, but is insoluble in hydrochloric acid. :
The precipitate redissolves in ammonium sulfide. ;
4. Reducing Agents.—If the solution of an alkali fuga is
treated with HCl and zinc, the tungstic acid at first precipitated by
the HCl is soon turned to a beautiful blue color, owing to the forma-
tion of W20s. ,
SnCle produces a yellow coloration at first, but on adding HCl .
and warming, a beautiful blue precipitate is bistied: This is one of -
the most sensitive reactions for tungstic acid. : a
5. Mercurous Nitrate precipitates white mercurous tungstate —
from neutral solutions.
6. Lead Acetate precipitates .white lead tungstate from neutral —
sclutions. 4

REACTIONS IN THE DRY WAY

The salt of phosphorus bead is colorless in the oxidizing flame,
and blue in the reducing flame, becoming blood red on the addition
of alittle FeSO. |

SELENIUM, Se. At. Wt. 79.2. Sp. Gr.=4.28-4.5. M. Pt.=217-220° C,

_Occurrence.—Although selenium is quite widely distributed in nature, —
it is invariably found in very small amounts, usually replacing sulfur, —
forming isomorphous compounds with lead, silver, copper, and mercury; —
clausthalite, PbSe; berzelianite, (CuAgTl)2Se; naumannite, (Age Pb)Se; —
tiemannite, HgSe; lehrbachite, (Pb,Hg)Se; onofrite, Hg(SeS); eu- —
eairite, (Ag,Cu)2Se. It is also found in small amounts in many varieties —
of pyrite and chalcopyrite, and indeed the small quantities which are —
found in these minerals form the chief source of the selenium of com- —
merce. By roasting these minerals (as in the manufacture of sulfuric —
acid) all of the selenium is volatilized, and is consequently deposited —
in the lead chambers as a mud from which it is extracted with a solution —
of potassium cyanide and afterwards precipitated with acid:

KCN+Se=KCNSe and KCNSe+HCl=HCN+KCI-+Se.

Selenium, like sulfur, exists in two allotropic forms. The modi- 4
fication soluble in carbon disulfide is obtained by reducing selenious acid —
in the cold with sulfurous acid; it is a brick-red powder. After heating ©

METALS OF THE H,S GROUP 497

this red selenium with hot water for some time, it is changed into black
selenium, and is then insoluble in carbon disulfide.

On heating in the air, selenium burns. with a bluish flame (giving
off an odor similar to that of rotten radishes) forming white, crystalline
selenium dioxide, SeO2, which will sublime on being heated in a stream
of oxygen. . Selenium forms one oxide, SeOz, and two acids: selenious
acid, H2SeOs, and selenic acid, H2SeOx.

Selenious acid, H2SeO3, is obtained in the form of long colorless
needles by oxidizing selenium with nitric acid or aqua regia,* or by
the solution of its anhydride, SeOe, in water. Unlike sulfurous acid,
it is not changed on standing in the air into selenic acid; but, on the
contrary, is reduced by dust, etc., to red selenium. The acid is dibasic,
and forms salts in which either one or both of the hydrogen atoms are
replaced by metal.

The acid salts are all soluble in water, but the neutral salts are all
insoluble with the exception of those of the alkalies.

Selenic acid, H2SeOz, is obtained in solution by conducting chlorine
into water which contains either suspended selenium or dissolved
selenious acid:

Se+3Cle+4H20 = H2Se04+6HCI.

Sodium seleniate is obtained by fusing selenium with sodium car-
bonate and potassium nitrate. Selenic acid is a dibasic acid and
behaves similarly to a peroxide, evolving chlorine when boiled with
concentrated hydrochloric acid, being reduced to selenious acid:

- HeSe04+2HCl=H.0+H28e03+Cl f.

REACTIONS IN THE WET: WAY

(a) Selenious Acid
A solution of either potassium selenite or of free selenious acid
should be used.

1. HeS produces a lemon-yellow precipitate, consisting of selenium
and sulfur, from solutions in water or in dilute hydrochloric acid:

H2SeO03 + 2H2S = 3H2O +Se +28.

The precipitate is soluble in ammonium sulfide.
2. Reducing Agents.
SOz2 precipitates red selenium. The solution may be hot or cold,

* If a solution of selenium in aqua regia is evaporated, considerable selenium
is lost by volatilization; the addition of KCl or NaCl to the solution prevents
such loss.

498 REACTIONS OF SOME OF THE RARER METALS ~

and contain little or much hydrochloric or sulfuric acid; by long boiling q ;
the precipitate turns grayish black. a
SnCly precipitates red selenium even in the presence of considell q
able sulfuric acid. 3
FeSO. immediately precipitates selenium from concentrated solu- —
tions of selenious acid containing hydrochloric acid; from dilute solu- —
tions the precipitation takes place very slowly and incompletely if
much sulfuric acid is present. |
Hydroxylamine hydrochloride precipitates selenium from solu- —
tions of selenious acid containing sulfuric or hydrochloric acid on long —
boiling; the precipitated selenium is red at first but finally becomes —
gray (difference from tellurium). 4
Hydrazine hydrochloride precipitates selenium from hot acid —
and alkaline solutions; the selenium is red at first and finally gray.
Hydriodic Acid (KI and HCl) Dt ag a red selenium in the 4
cold (difference from tellurium). %
Zinc precipitates red selenium from acid solutions: the zine becomes —
coated with Se and looks as if covered with Cu. 4
3. BaCle precipitates from neutral solutions white barium selenite,
BaSeOs, soluble in dilute acids. a
4. CuSO. produces a greenish-blue, crystalline precipitate (differ- 4
ence from selenic acid).

(b) Selenic Acid

1. HeS causes no precipitation unless the solution is boiled with
hydrochloric acid. In the latter case the selenic acid is reduced first to
selenious acid and then to selenium, which precipitates together with .
free sulfur.

2. BaCl2 gives a white precipitate of barium seleniate, BaSeQOu,
insoluble in water and in dilute acids, soluble, with evolution of chlorine,

|
A solution of potassium seleniate should be used.
on being boiled with hydrochloric acid:

BaSe0.4+4HCl=BaCle+ H2Se03+Cle fT +H20.

3. CuSO, produces no precipitation.
4. SO does not reduce selenic acid except by long boiling with —
hydrochloric acid. The reduction takes places more readily with —
hydrazine.

METALS OF THE HS GROUP 499

Method for Testing Sulfuric Acid for Selenium *

Add 5 or 6 drops of the acid to be tested to a freshly prepared solution
of a little codein in sulfuric acid; if selenium is present, a green coloration will
be apparent. The test is a very delicate one.

REACTIONS IN THE DRY WAY

All selenium compounds emit the odor of decayed radishes on being
mixed with sodium carbonate and heated on charcoal before the blow-
pipe.

If a selenium compound is heated at the end of a thread of asbes-
tos in the upper reducing flame of the Bunsen burner, it will be reduced
to selenium; and if a test-tube filled with water is held above the flame,
a red coating of selenium will be deposited upon the glass.t If a few
drops of concentrated sulfuric acid { are placed in a larger test-tube
(large enough to hold the smaller test-tube) and the tube on which the
selenium is deposited is emptied and placed within the’ larger tube,
the selenium will dissolve § in the sulfuric acid, forming a green solu-
tion; but on the addition of water, red selenium will be reprecipitated
(difference from tellurium) :

SeSO3-+ H20 =Se+He2S0O..

Green.

TELLURIUM, Te. At. Wt. 127.5. Sp.Gr.=6.1-6.4. M. Pt.=452° C.

Occurrence.—Tellurium is a rarer element than selenium, always
occurring in. the form of a telluride, and usually combined with the
noble metals: calaverite, (Au,Ag)Te2; krennerite, (Au,Ag)Te2; sylvan-
ite, (Au,Ag)Tes4; nagyagite, AugzSbePbioTesSi5; coloradoite, HgTe;
silver telluride, AgeTe; and often in small amounts in galena and copper
ores. Emmonsite of Cripple Creek, Colorado, is a ferric telluride with
‘70.71 per cent TeOo and 22.76 per cent Fe2O3. Tellurium itself is
a bluish-white, brittle substance, which can be distilled in a stream of
hydrogen. It burns in the air with a bluish-green flame, forming
tellurium dioxide, TeOz. It is insoluble in carbon disulfide, and can

. DRAGENDORIF, Chem. Zentralb., 1900, 944.

+ Cf. p. 68. :

t The sulfuric acid should be freed from water by heating in a platinum crucible
to a temperature just below the boiling-point, and the crucible hin its contents
allowed to cool in a desiccator.

§ Slowly in the cold, readily on warming.

500 REACTIONS OF SOME OF THE RARER METALS

be oxidized by means of nitric acid to tellurous acid. On being fused
with potassium cyanide, out of contact with the air, it is changed to
potassium telluride,

2KCN +Te = KoTe+(CN)s2,

which dissolves in water, forming a cherry-red solution. If air is con-
ducted through this solution, the tellurium is. precipitated in the form
of a black powder (difference from selenium) :

KeTe+H20+0=2KOH+Te.

Tellurium may be separated from selenium by means of this last
reaction. The two metals are fused with potassium cyanide, the melt
is treated with water, and the tellurium precipitated by passing a current
of air through the solution; the selenium is precipitated from the filtrate
by acidifying with hydrochloric acid. Tellurium forms two oxides:
TeOz and TeQs3.

Tellurium dioxide (the anhydride of tellurous acid) is usually obtained in
the form of a white mass, which melts on gentle heating, forming a yellow
liquid. Tellurium dioxide does not sublime (difference from selenium). It is
scarcely soluble at all in water, is slightly soluble in ammonia and in dilute
acids, but readily soluble in concentrated acids or in caustic potash solutions.
TeO, dissolves in fairly concentrated sulfuric acid, forming the basic sulfate,
Te.0;-SO., while with nitric acid it forms the basic nitrate, TexO;(OH)NO;.
Both of these compounds are hydrolized readily, forming insoluble tellurous
acid; the latter in turn loses water and forms the anhydride.

On dissolving TeO. in caustic potash, potassium tellurite, K,TeO;, is
obtained. Only the alkali tellurites are soluble in water.

Tellurium trioxide (telluric anhydride) is formed by heating tellurie acid.
It is a yellow powder, insoluble in water and nitric acids, scarcely affected by
boiling with concentrated hydrochloric acid, but is readily dissolved by boiling
with a concentrated solution..of potassium hydroxide (but not by sodium
hydroxide), forming potassium tellurate.

Telluric acid, H,.TeO.,+2H;0, is a very weak acid, obtained by oxidizing
tellurous acid with chromic acid, and precipitating the telluric acid by the
addition of concentrated nitric acid. The acid forms a colorless crystalline
mass, is readily soluble in water, and is converted by means of concentrated
hydrochloric acid into tellurous acid, with evolution of chlorine. The acid
dissolves readily in caustic potash (or soda) solution, forming the readily
soluble alkali tellurate, which reacts strongly alkaline in aqueous solution.

By gently heating the hydrated telluric acid, the anhydrous acid, H:TeQ,,
is obtained in the form of a white powder and is totally different from the

hydrated acid. The latter is soluble in water and in caustic alkalies, and is ~
completely reduced by boiling with concentrated hydrochloric acid; but the —

anhydrous acid is insoluble in water and in concentrated sodium hydroxide
solution, and is only very slightly attacked by boiling, concentrated hydro-

vr ee

_—_— wees Se 5 hattatlmn, — =

7 , oe

Es ee ee a are eer nS) aS ee ee ee A

ES a

METALS OF THE H,S GROUP 501

chloric acid, although readily soluble in warm potassium hydroxide solu-
tion.

Only the alkali tellurates are soluble in water; the others are usually
obtained in the form of amorphous precipitates soluble in acids.

REACTIONS IN THE WET WAY
(a) Tellurous Acid
A solution of potassium tellurite, K,TeO;, should be used.

1. HeS precipitates from acid solutions brown TeS2, which is readily
soluble in ammonium sulfide.

2. Reducing Agents.

SO» precipitates tellurium completely from dilute hydrochloric
acid solutions in the form of a black powder, even in the presence of
tartaric acid; but from a solution containing considerable hydro-
chloric acid no tellurium is precipitated even opr boiling (difference
from selenium). The separation of the selenium from tellurium
can be accomplished in hydrochloric acid, sp. gr. 1.18.

SnCl2 or Zn causes black tellurium to precipitate from solutions
tvhich are not too acid.

H3PO3 precipitates the tellurium only from concentrated solu-
tions, not at all from cold dilute solutions.

FeSO, reduces neither tellurous nor telluric acids (difference from
selenium).

Hydroxylamine Hydrochloride produces no precipitate in solutions
of tellurous acid containing mineral acids, but precipitates tellurium
completely by boiling the ammoniacal solutions for a long time.

INH.OH+Te037+2H*t > 4H20+N20 T +Te.

Hydrazine Hydrochloride precipitates black tellurium both from
acid and ammoniacal solutions:

NeoH4:-2HCI+TeO3 — 3H20+ Ne T +Te+2CI.

Hydriodic Acid (KI and HCl) produces no precipitation, but merely
a reddish-brown coloration that turns light yellow on boiling (differ-
ence from selenium).

Zinc (or Fe, Sb, Sn, Cd, Hg, Pb, Cu, etc.) COANE black tellu-

* yium.

3. HCl produces a white euscinabata of HeTeQ3.

502 REACTIONS OF SOME OF THE RARER METALS

(b) Telluric Acid.

A solution of potassium tellurate should be used.

1. HCl causes no precipitation; but if the solution is bailed chlorine
is evolved, and on dilution with water tellurous acid is precipitated. |

2. HeS and reducing agents have the same effect upon hot solu-
tions of tellurates as upon tellurites.

3. Lead Salts precipitate difficultly soluble lead tellurate.

4. Hydrazine Hydrochloride precipitates all the tellurium as a
black powder by long boiling of the acid or alkaline solution:

3(N2H4:-2HCl) +2Te04" — 8H20+2H*+6Cl-+3Ne2 Tf +2Te.

Detection of Selenium and Tellurium in Ores

Principle-—The finely powdered, dry ore is heated in a current of chlorine;
the chlorides of sulfur, selenium, arsenic, antimony and iron are volatilized
and may be absorbed in dilute hydrochloric acid.

Procedure.—Place the finely powdered, dry mineral in a porcelain boat,
push this into a tube of difficultly fusible glass and connect the tube on one
side with the wash-bottle and on the other side with a 10-bulb Meyer tube.

Prepare chlorine gas from a Kipp generator containing chloride of lime
and hydrochloric acid, wash the gas by passing it through a bottle contain-
ing water and dry it by passing it through concentrated sulfuric acid. Pass
the chlorine gas through the tube and when all the air has been expelled,
begin heating the substance, at first very gently. Fumes quickly begin to
form, showing that the action of chlorine upon the substance has started.
Soon vapors of sulfur chloride, 8.Cl, (B. P. 64°) begin to condense in the
front end of the tube in the form of drops; drive these over into the receiver
by carefully heating the tube. As the temperature is raised a little, a white
sublimate of selenium chloride, SeCl, (B. P. 200°) forms in the front end of
the tube; drive this over into the receiver in the same way. Now heat the
boat hotter and soon brown vapors of ferric chloride will be evolved, some of
these will condense, forming glistening scales which must also be driven over
into the receiver. Continue heating until finally no more vapors are evolved.
Transfer the contents of the 10-bulb tube to a porcelain dish, add 0.5 gm. of
potassium chloride to prevent loss of selenium chloride by volatilization, and
evaporate the solution to dryness on the water-bath. Dissolve: the residue in
as little hydrochloric acid as possible and treat the solution with stannous
chloride. If tellurium is present a black precipitate is obtained which may
also contain selenium. Filter through an asbestos filter, wash the residue with
dilute hydrochloric acid, place the filter and asbestos in a small test-tube, boil
it with concentrated hydrochloric acid until no more black spots are visible in
the asbestos, dilute with water and filter. LEvaporate this filtrate to dryness
on the water-bath, dissolve the residue in 10 ce. of HCl, sp. gr. 1.175, and pass —
SO, gas into the hot solution. Filter off any precipitate of red selenium that
may form, Dilute the filtrate with considerable water and again pass SO, into

METALS OF THE H,S GROUP 503

the hot solution, which should precipitate the tellurium as a black powder.
Identify the selenium and the tellurium by the sail reactions given on pp. 499
and below.

REACTIONS IN THE DRY WAY

Metallic tellurium is formed by heating any telluride in the upper
reducing flame, and can be collected on the lower surface of a test-tube,
which is filled with water, in the form of a black film, soluble in
concentrated sulfuric acid. The latter solution is of a carmine-red
color (difference from selenium); on the addition of water black
tellurium is deposited:

TeSO3 +H:0O =Te -|- He2SO..

Carmine red

THE PLATINUM METALS

PLATINUM, PALLADIUM, RHODIUM, OSMIUM, RUTHENIUM AND 4
IRIDIUM ;

Platinum has been described a'ready on p. 266.

PALLADIUM, Pd. At. Wt. 106.7. Sp. Gr.=11.8. M. Pt.=1550° C.

Occurrence.—The platinum metals form an isodimorphous group, — |

but only in the case of palladium are both forms known—the isometric
and the hexagonal:

(a) Isometric System. (b) Hexagonal System.
Platinum (Pt, Fe). Tridosmium (Sysserskit) (Ir, Os). %
Iridium (Ir, Pt). Osmiridium (Newjanskit) (Ir, Os,
Platinum iridium (Pt, Ir, Rh). Pt, Rh, Ru) or (Os, Ir, Rh).
Palladium (Pd, Pt, Ir). Palladium (Pd, Pt, Ir).

Properties—Rolled, hammered, or cast palladium possesses an
almost silver-white color, but when precipitated from solutions it is
in the form of a black powder. If it is suspended in water when in
the finely divided form, it is transparent with a reddish color. Palla-
dium has the lowest melting-point of all the platinum metals, viz.,
1550° C. On being heated in the air, it appears bluish, owing to the
formation of PdeO; the latter, however, is decomposed by stronger —
heating.

Behavior towards acids: Although the other platinum metals are
attacked by no acid except aqua regia, palladium is dissolved slowly

by warm nitric acid (also in the cold when it is alloyed with other —

metals such as Cu, Ag, etc.), forming a brown solution of Pd(NQs3)o.

Finely divided, precipitated palladium is soluble in hydrochloric
acid wher exposed to the action of air at the same time, and less readily
soluble in sulfuric acid. It is readily attacked by fusing with potas-
sium pyrosulfate, forming soluble palladium sulfate, PdSOz. —

The best solvent for palladium is aqua regia.

Finely divided palladium has the very characteristic property of
being able to absorb almost 700 times its own volume of hydrogen,
and possesses consequently a very strong catalytic action. If hydro-

504

THE PLATINUM METALS 505

gen and oxygen (air) are conducted at the same time over some gently
ignited, finely divided, metallic palladium, the hydrogen is burnt to
water, and in the same way carbon monoxide may be changed to carbon
dioxide. Methane, however, is only decomposed by igniting the
palladium more strongly, so that this gives us a method for separa-
ting methane from a mixture of H and CO (cf. Vol. 2, Gas Analysis).

Colloidal palladium preparations show marked catalytic effects.
Thus by passing hydrogen gas into solutions of unsaturated crganic
compounds a direct reduction (hydrogenation) often takes place if
a little colloidal palladium is present.

Palladium forms two oxides, both of which posses$ strongly basic
properties: PdO and PdOzg. From the former the palladous, and from
the latter the palladic, compounds are derived. The palladous com-
pounds are much more stable than the palladic compounds, and the
latter constantly exhibit the tendency to change into the former.

By dissolving finely divided palladium in hydrochloric acid, palla-
dous chloride is formed; or, better, by dissolving the metal in aqua
regia, in which case a mixture of palladous and palladic chlorides is
at first obtained. If this solution, however, is evaporated to dryness,
palladic chloride loses chlorine and is completely changed into palla-
dous chloride, so that on treating the residue with water a solution of
palladous chloride is obtained. Since palladic chloride is decomposed
completely by evaporation, it is bday that palladic chloride cannot
exist in hot solutions.

REACTIONS IN THE WET WAY

(a) Palladous Compounds
Use a solution of palladous chloride, PdCh.

1. H2S precipitates black palladous sulfide from acid and neutral
solutions. The precipitate, PdS, is insoluble in ammonium sulfide,
but soluble in boiling hydrochloric acid, or more readily in aqua regia.

2. KOH or NaOH precipitates a brown basic salt, soluble in an
excess of the reagent. If the solution is acidified with HCl, then
KOH produces no precipitate (difference from platinum).

3. NaCOz produces a brown precipitate of palladous hydroxide,
Pd(OH)s, soluble in excess but reprecipitated on boiling.

4. NH4OH gives a flesh-colored precipitate of [Pd(NH3)2Cle], *
soluble in an excess of ammonia, forming a colorless solution (eon-
taining palladodiamine chloride, Pd(NH3)4Cle), from which yellow

* This compound is an isomer of palladosamine chloride, and is often written
thus: PdCh,Pd (N Hs) aCle.

506 REACTIONS OF SOME OF THE RARER METALS |

crystalline palladosamine chloride, Pd(NH3)2Cle, is precipitated on the
addition of hydrochloric acid. The latter compound is difficultly — a
soluble in dilute hydrochloric acid and is used for the preparation of —
pure palladium. pris

In a solution of palladous nitrate, ammonia causes no precipitation,
but forms colorless palladodiamine nitrate, Pd(NH3)4(NO3)2.

5. NH,Cl. If a solution of palladous chloride or of sodium-palla-
dous chloride is treated with ammonium chloride and evaporated to
dryness on the water-bath, the residue is soluble in a very little water.
If the solution is acidified with nitric acid, gradually all the palladium
is precipitated as red (NH4)2[PdCle] (difference from platinum).

6. KCl, when added to a concentrated solution, causes the pre-

cipitation of difficultly soluble, reddish-brown Ke{[PdCl4] (octahedrons). C

7. HI or KI produces a black precipitate of palladous iodide, even
in very dilute solutions. The precipitate is insoluble in water, alcohol,
ether, and HI, but soluble in KI and NH3. (This and the following
reaction are characteristic of palladium.) .

8. Hg(CN)2 produces a _ yellowish-white gelatinous precipitate
of palladous cyanide, Pd(CN)e, difficultly soluble in HCl, readily
soluble in KCN and NH3. On being ignited, the spongy metal remains.

9. Nitroso-6-naphthol (a saturated solution in 50 per cent acetic
acid) gives a voluminous, brown precipitate of Pd(CiopHegNO2)2 even
in the most dilute solutions (difference from platinum).

10. Reducing Agents.—H2SOs, formic acid, HCOOH, Zn, Fe,
FeSO4, CueCle,* alcohol, and CO7f reduce palladium salts to the
metal itself.

In the presence of HCl, stannous chloride forms at first a red, then
a brown, and finally a green, solution; but in the absence of the acid,
SnCle causes a partial reduction to the metal and the solution turns
green.

(b) Palladic Compounds

These give the same reactions as palladous compounds, on account
of their being readily changed into the latter. The principal differ-
ence, however, is the insolubility of the ammonium salt of chloropalladic

* In the presence of considerable NaCl or HCl there is no reduction with CusCh.
+ PdCl,+CO+H,0 =2HCl1+CO;+Pd. This reaction enables one to detect
small amounts of CO in gas mixtures; e.g., in the air. For this purpose the gas
is led through a narrow glass tube into 10 cc. of a solution which contains 1 mgm.
of PdCl, and 2 drops of dilute HCl. If CO is present, black Pd will be deposited,
and the solution will become decolorized little by little. (Porarin and Drourn,

Compt. rend., 126, 938.) If too much HCl is present the reduction will not take __

place unless NaC,H;O, is added.

THE PLATINUM METALS 507

acid. If a concentrated, cold solution of palladous chloride is shaken
with chlorine water and then treated with ammonium chloride, a red
crystalline precipitate of (NH4)2[PdClg] is soon formed.

REACTIONS IN THE DRY WAY

All palladium compounds are decomposed on ignition, leaving
behind the metal, which is soluble in nitric acid or in aqua regia, and
_the solution thus obtained can be tested by the above reactions.

RHODIUM, Rh. At. Wt. 103.0. Sp.Gr.=12.6. M. Pt.=1920°?

Properties.— Rhodium possesses the color and luster of aluminium; it
is more infusible than platinum and melts at about 1920°; on cooling
the hot metal it sputters and appears bluish, owing to oxidation. The
solubility of rhodium depends entirely upon the fineness of the material.

When precipitated from a solution of its chloride by means of
formic acid or other reducing agents at a temperature not exceeding
100°, it exists in an extremely finely divided state (rhodium black)
and dissolves readily in boiling, concentrated sulfuric acid, or more
readily in aqua regia. If, however, the finely divided metal is ignited
strongly, it becomes (like the compact metal) almost insoluble in aqua
regia. |

If rhodium is alloyed with large amounts of other metals (Pb, Zn,
Bi, Cu, etc.), it is left in a finely divided condition after treatment of
the alloy with acids, and is consequently soluble in-aqua regia. When
it is alloyed with much platinum or palladium a considerable amount
of rhodium will dissolve in aqua regia; but when it is alloyed with a
little platinum, most of the rhodium and a part of the platinum remain
undissolved.

On being fused with potassium pyrosulfate, potassium rhodium
sulfate is formed, which dissolves in water, forming a yellow solution,
but becomes red on the addition of HCl. :

Rhodium forms three oxides: RhO, RhzO3, and RhOg; all possess
a well-defined basic nature. The sesquioxide, Rh2O3, alone * forms a
series of salts, of which sodium-rhodium chloride is the most important
for the analytical chemist; when in this form it is easiest to bring
rhodium into solution. This salt is prepared by mixing the finely
divided metal very intimately with twice as much dry sodium chloride,

* A sodium-rhodium sulfite of the formula 4Rh(SO3),6Na:sS0;+9H20 was
prepared by Bunsen.

508 REACTIONS OF SOME OF THE RARER METALS

placing it in a porcelain boat and gently igniting it in a current of
moist chlorine gas. The salt thus formed has the composition
Nag{RhCl¢] and is soluble in water (45 parts of water dissolve 1 part
of the salt). From this solution large, dark-red, glistening triclinic
prisms of Na3[RhCle6]+9H20 can be crystallized out.

REACTIONS IN THE WET WAY
Use a solution of sodium rhodium chloride Na;[RhClk.]

1. H2S precipitates (very slowly in the cold, but much more quickly-
on warming) black rhodium sulfide, Rh2S3; insoluble in (NH4)28,
soluble in nitric acid.

2. KOH and NaOH produce at first no precipitate; but after —
standing some time a yellow precipitate of rhodium hydroxide,
Rh(OH)3+H20, separates out. The precipitate is soluble in an excess
of the reagent, but it is reprecipitated on boiling in the form of brown-
ish-black Rh(OH)3.

In a solution of potassium rhodium sulfate, KOH precipitates the |
yellow compound immediately.

On adding KOH to a solution of rhodium chloride, at first no
precipitate is produced; but on the addition of a little alcohol brown-
ish-black rhodium hydroxide is deposited.

3. NH4OH produces (in concentrated solutions and after stand-
ing some time) a yellow precipitate of chlorpurpureorhodium chloride,
Rh(NHs3)5Cls, insoluble in hydrochloric acid.

4. KNOs, on being warmed with sodium rhodium chloride solu-
tion, causes the precipitation of difficultly soluble, orange-yellow
Ks[Rh(NO2)¢6], soluble in HCl.

5. Reducing Agents.—Formic acid in the presence of ammonium
acetate, precipitates the black metal, as does zinc in the presence of
acids.

REACTIONS IN THE DRY WAY

All rhodium compounds are reduced to metal on being heated in
a stream of hydrogen, or by heating on charcoal with sodium car-
bonate before the blowpipe. The metal is easily recognized by its
insolubility in aqua regia, its being brought into solution by fusing
with potassium pyrosulfate and then treating with water, and by the
formation of the brown hydroxide when.KOH and a little alcohol are
added to the solution thus obtained.

THE PLATINUM METALS 509

OSMIUM, Os. At. Wt. 190.9. Sp. Gr.=21.3-22.48. M. Pt.=about
2700°

Osmium and ruthenium are distinguished from the other platinum
metals by their forming volatile oxides.

Properties—The compact metal possesses a bluish-white color, very
similar to zine, and is the heaviest of all metals. It can be melted by
heating in an electric furnace.* Very finely divided osmium is oxidized
by the air + at ordinary temperatures, and at about 400° C. it ignites
and burns rapidly to OsO4, which is volatile at 100° C. The denser
the metal the higher the temperature necessary to effect the oxidation.

Behavior towards Acids—In the compact condition osmium is in-
soluble in all acids; but in the finely divided state (as obtained by
treating its zinc alloy with nitric acid) it is soluble in nitric acid, more
soluble in aqua regia, and most soluble in fuming nitric acid, forming
osmium tetroxide; the latter can be separated from the solution by
distillation.

Compact osmium is brought into solution by fusing with NaOH and
either KNOs or KClO3. The melt contains a salt of perosmic acid

(OsO4).

Osmium forms five oxides:

OsO Os20sz, OsOg,

Osmious oxide, grayish-black, Osmium sesquioxide, black, Osmic oxide, black-gray,

insoluble in acids insoluble in acids insoluble in acids

[OsOs], OsO.
Osmic acid, known only in Perosmic acid, colorless needles
derivatives soluble in water

Osmium tetroxide, OsO4, (the anhydride of perosmic acid,) is the
most important osmium compound in the eyes of the analytical chemist,
and is obtained by the oxidation of the substance in the air by dis-
solving the finely divided metal in fuming nitric acid or in aqua regia,
or by fusing with NaOH and KNOgz or KCI1Os, treating the melt with
nitric acid and distilling. Osmium tetroxide is a colorless, crystalline
mass which sublimes at a comparatively low temperature and melts,
forming colorless vapors at 100° C. The vapor has a chlorine-like
odor, attacks the mucous membrane, and is poisonous.

The chlorides of osmium can be obtained only in the dry way;
OsCle, OsCls, OsCla are known. ‘The potassium salt of the hypothet-
ical hydrochlorosmic acid, H2OsCle, forms dark-red octahedrons, soluble

*F. Myuius and R. Dintz, Ber., 1898, 3187.
7 Cf. Or. Suc, Z. anorgan. Chemie, 19, 332.

510 REACTIONS OF SOME OF THE RARER METALS

in water and decomposed by boiling the solution. By heating finely — a
divided osmium with KCl in a current of chlorine, K2[OsCle] is formed; — ;
~ it dissolves in cold water, forming a red solution.

REACTIONS IN THE WET WAY
Use a solution of K,[OsCl,].

1. If a solution of osmium chloride is treated with dilute nitric
acid, the mixture distilled from a small retort, and the vapors received —
in caustic soda solution, the latter will be colored yellow, owing to the
formation of potassium osmiate. If this solution is now acidified, —
osmium tetroxide is set free, and can be recognized by its very penetra-
ting odor. On adding a little sodium thiosulfate to the acid solution
and warming, a brown precipitate of osmium sulfide is formed. ‘

2. HeS precipitates brownish-black osmium sulfide, insoluble
in ammonium sulfide. ay ‘

3. KOH, NH4sOH or K2COz precipitate reddish-brown osmium _
hydroxide Os(OH)a.

4. Reducing Agents.—If the solution of the chloride is treated
with tannic acid and alcohol and a little hydrochloric acid is added, —
it is colored dark-blue, owing to the formation of osmium dichloride, —
OsCle; KI colors the solution a deep reddish-purple.

5. Indigo is decolorized by solutions containing OsO4. Ferrous
sulfate precipitates black osmium dioxide; stannous chloride produces
a brown precipitate soluble in HCl, forming a brown solution.

REACTIONS IN THE DRY WAY

All osmium compounds are reduced to metal on being heated in
a stream of hydrogen.

RUTHENIUM, Ru. At. Wt. 101.7: Sp. Gr.=12.26, crystallized; 11.0,
fused. M. Pt.=above 1950°

Properties.—Ruthenium exists in the form of a dark gray or black
powder, and in the form of bright porous sticks; it is brittle, can be
powdered, and is melted in the oxyhydrogen flame.

On being melted, a part of the ruthenium is oxidized to ruthenium
tetroxide, a volatile substance having a penetrating odor similar to
that of OsO4. The molten metal spurts on cooling. .

Behavior towards Acids—Ruthenium is almost completely insoluble —
in all acids, even aqua regia. By fusing with KOH and KNOg3 (or _
KCIO3) it is oxidized to potassium rutheniate, K2[/RuQO,]. a

THE PLATINUM METALS oll

On heating with NaCl in a current of chlorine, soluble Ke[RuCle]
is formed. The solution in water of the greenish-black melt is of an
orange-yellow color, and colors the human skin black. Ruthenium is
unaffected by fusing with potassium pyrosulfate.

If forms the following oxides:

RuO, RuzOz, RuOe, [RuOs], [RueO7z], RuOs.

The most important of the oxides is RuO4. It is formed:

(a) By roasting the metal itself, or its oxide, above 1000° C. (osmium
forms the volatile tetroxide at 400° C.).

(b) By fusing the metal with KOH and KNOs in a silver crucible, |
dissolving the melt in water, saturating the cold solution with chlorine
gas, and distilling the solution from a small retort:

KeRu04+Cle = 2KCl1+RuO..

(d) By distilling potassium-ruthenium chloride with KClO3 and
HCl.

By distillling a dilute solution after the addition of nitric acid, ne RuO4
will be evolved* (difference from osmium),

Ruthenium tetroxide forms gold-yellow, glistening, orthorhombic
needles that are very volatile and emit a characteristic odor; it boils
at 100° C. and is only slightly soluble in water. It is changed by the
addition of alcohol and HCl into ruthenium trichloride, RuCls (or
sesquichloride, RugCle). If the solution of the latter salt is made
ammoniacal, treated with sodium thiosulfate and warmed, an intense
reddish-violet coloration will be produced. (This is a very sensitive
and characteristic reaction.)

On treating a solution of potassium ‘rutheniate with nitric acid,
black Ru(OH)s3 is precipitated; it dissolves in hydrochloric acid, form-
ing a yellow solution of RuCls. .

REACTIONS IN THE WET WAY

Use a solution of RuCl,.

1. H2S produces no precipitation at first, but after some time the
solution becomes azure-blue, and brown ruthenium sulfide is precipi-
tated (very sensitive and characteristic).

* In the presence of concentrated nitric acid, however, RuO, is formed:

2K,Ru0O,+4HNO; =4KNO;+Ru(OH).,+ RuQ,.

512 REACTIONS OF SOME OF THE RARER METALS

2. (NH,)2S precipitates the brownish-black sulfide, difficultly —

soluble in an excess of the reagent.

3. KOH and NaOH precipitate black ruthenium hydroxide, ;

Ru(OH)s, soluble in acids but insoluble in alkalies.
4. KCNS, in the absence of other platinum metals, produces

gradually a red, then a purple, and on warming a violet, coloration (very _

characteristic).

5. KNO: imparts an orange-yellow color to the solution, owing
to the formation of Ks{Ru(NO2).6] becoming a beautiful dark red on
the addition of a little colorless. ammonium sulfide; on adding more
ammonium sulfide, brown ruthenium sulfide is precipitated.

6. Zinc at first colors the solution of the chloride azure blue, but —

subsequently the solution is decolorized and ruthenium itself is pre-
cipitated.

7. Hydroxylamine reduces ruthenium tetrachloride to ruthenium |

trichloride (difference from platinum),

IRIDIUM, Ir. At. Wt. 193.1. Sp. Gr. 22.4. M. Pt.=2300°?

Properties —When produced by the ignition of iridium ammonium
chloride, it is obtained in the form of a gray, spongy mass, very dif-
ficultly soluble in aqua regia. After being strongly ignited, it is almost
completely insoluble in aqua regia. |

It is more soluble in aqua regia after it has been precipitated from |

solutions in a very finely divided form by means of formic acid, or
when it is alloyed with other metals (Au, Ag). The metal is unaffected
by fusing with potassium pyrosulfate (difference from rhodium). It is
oxodized by fusing with NaOH and KNOs in a silver crucible, but the
compound formed (Ir2O3 combined with sodium) is only partly soluble
in water. If the melt is treated with aqua regia, however, a dark-red
solution of Nag[IrCl's will be obtained.

By heating the metal with NaCl in a current of chlorine, Nag[IrCle]
is readily obtained.

Iridium forms the following oxides:

Ir203,? IrOz, and the hydroxide Ir(OH)a.
Bluish black Needles with a Indigo-blue
metallic luster powder

The dark color of the chlorides is very characteristic:
IrCl3,* IrCh,

Dark green and brown : Black

* According to W. Patmagr, Z. anorg. Chem., 10, 322-326 (1895), IrCls and

its double salts exist in two modifications: dark green nied brown.

. - ‘ &
ee Eee ee ae

THE PLATINUM METALS 613

REACTIONS IN THE WET WAY
Use a solution of Na[IrC],].

1. HeS at first decolorizes the solution, owing to the reduction
of the tetrachloride to the trichloride, accompanied by the deposition
of sulfur; subsequently brown Ir2S3 is precipitated, readily soluble
in (NH4)28.

2. (NH4)2S precipitates the same compound.

3. NaOH, on being added to the solution, changes the color from
dark red to green; on warming the solution it is at first colored reddish
and finally azure blue:

2IrCl4a+2Na0H = 2IrCI3+ NaCl+H20+Na0Cl.

If the solution is now acidified with HCl, a little alcohol * added
and then some KCl, there will be no precipitation, because the Ks[IrClg|
formed is readily soluble in water and in KCI solution (difference from
platinum).

4. KCl precipitates brownish-black potassium iridium chloride,
Ko[IrCle], insoluble in KCl and in alcohol, difficultly soluble in water.

5. NH,Cl precipitates dark-red ammonium iridium chloride,
(NH4)2[IrCle], insoluble in a saturated solution of NH4Cl.

6. Reducing Agents usually change the solution to a greenish
color, owing to the reduction of the tetrachloride to trichloride; or the
solution is decolorized and the black, finely divided metal is deposited.

Thus if the solution is warmed with KNO,, an olive-green coloration is

produced:
IrCl,+KNO, =IrCl;+KCI+NO, 1.

If the solution is boiled for some time with an excess of KNOs, it becomes
yellow, and a part of the iridium separates out in the form of a yellowish-white
precipitate, difficultly soluble in cold hydrochloric acid or in boiling water.
The precipitate has the following composition: 3K;[Ir(NOz).|-K{IrCl.].

Oxalic acid, ferrous sulfate, stannous chloride, and hydroxylamine reduce
the tetrachloride to trichloride. Zinc reduces it to metal, and so does formic
acid on warming in the presence of ammonium acetate. If considerable
mineral acid is present the reduction takes place less readily.

7. Chlorine. If chlorine is conducted into a dilute solution of
iridium tetrachloride, the latter becomes reddish-violet according to
W. Palmaer,{ after some time the red color disappears and the solution

turns brown. The same red color has been observed by Foerster ¢

* The alcohol reduces NaOCl to NaCl.
+ Z. anorg. Chem., 10 (1895), 358.
t Z. Electrochem., 10, 715,

REACTIONS OF SOME OF THE RARER METALS:

514

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up yy ‘ssauhsp 0} fymojs worynjos ay} aynsodvagy | autumnpfixotphy ppv ‘4az0n joy ur ayoprdrdead ayy aajossigy | -YNSa4 ay?
8a) pun
) ; ‘sa27nm | 8O Jo 90u9
PIOUWEN “TOPd*eN] Nowatog ‘POMUVHN) “LOAPCHN) “10%d*(HN)] Gtvuratoayg | 977727 » |- 892d ayy
uranprsas | SMOUS Spy
. van “pautof sr aypdroesd » sp buo) sp woyngos ay} 07 apitozyo wnruoMuD ppbv puv ft JayoM ynM anprsa4 ayy aajos |~[2S unTUso
9Y} 7041} bux009 Jajfy *(7JaUL JOU JSNUL SSD AY) WAL}S AULLO]YI JSOU D UL SsaUpas ]]Np 0} JooY puD YDS -sep ‘sayja | FO eyeqtdto
j | wommos YM anprsau fap ay) Tim UayT, “prop avuLtof fo wor ppD ay} Uo dnoLb sry? YR payopdrdard uaaq aany | y 2 fo -o1d WMOIq
g | how yoryn pojaw ubrasof tayjo 10 “ugy fun aaowas 0) Lapso Ut WOH 104 yun ‘urwby pj0o Buruorag «azfo ‘sypjau aouodvas | V° “uoumpos
2 | buruwmas ay) prog *burynoa fivsb v fo wsof ay} ur aqn) ay) fo suorj1od 1apjoo ay} UO pasuapuUod fiqyuanbasqns st | -quaso i! ay} yay pun
a iL ‘pazyynjoa aq 07 bY fun sasnva yoryn ‘uaborphy fo woos D U2 7o0g ULy/ad10d Db UA ayoyrdioaid partp ay) 3a] S{qeqoad "O%S*D NT
i | | st ony | 2 ? PPV
“NOLLO'I0g ‘(3H ‘ny “MY “IT ‘Pd “Iq) ALVLIaloaug ‘moped st |*OSO fo topo
uorjznjos |? %? 970%
“La7OM Pun aD}aID WNAUOWUD YR ayoRLd2I—Ld BY} YSDN PUY Jayry “lasuapuod moy-uinjas | eoroyyo ||P“? 10H
D YIM payvauuos ysoyf ay? yr ‘sinoy Yo1adas Lof WoLNjos ay} 1109 PUD pLID IwLof pun aynjazv wnruoMmun ppy|ayr yy | Yer Afyproo
asDo Sty? UT
‘guosoid
*(S[VJOUL SUTUTBUIAL oY} SUIB}JUOD) NOLLATOg snoTaAdy ‘NOILQTOg IvaugHiq | 9q ABU
SQ. ‘moroA
"s89].10j09 sinaddy jov.ujxa ay} ay} YUN sayja Ynn fypayadas uoynjos |sawo0de q
ay? ayoygy “jauunf fuojosndas v fo suvawu fq sahn sayja ay} aaowas pun ‘pinby ay} aynys ‘uayja fo ‘90 maf v ppp ‘Buyoos wafy | HORN O4t IT
‘ELV TTILSICY

"S[BIOUL SUTUTSUTAL oY} SUTB}OD 4.10301
NH 97NQUp YUN 701) “7L0}aL PJDWS DY Ur

UoYNjos ay) BOD} “Saplloyyo SB uOTyNIOS ul yUaserd aq 0} powmnsse ore ‘AIndJouT pu Pjo3 Jo oouesaid oy} ur ‘speyou WNUTZeT xIS OTT,

« SIVIGIN WONILWId HHL AO SISA TVYNYV

THE PLATINUM METALS

515

“ATWO UOI}VUUTEXE 9a27D;DNH OYZ JOY POYJOU SIY} PUSUTUIODAL AOU, “ZEIE “EST ‘~9gG ‘ZLaiq pu sAriA, x

"**ONM pus

GHA Buisny 190}J8 NY+ I] JO} poururexa st 41 ‘soygns yOu seOp SI} Jt ‘poysoder ssooord oY} Puv YO polsz]y St 41 ‘10jBM YIM BurjoOwI}xXS 1ozJe SuTeUIe! ONpisal B JT

*10[09 MOTIEA ¥B [HAJ PUB Joye OY} 07 JO[OD oON{q 10 MOTE B JaBduUTT [[IM plow ormOIgD

‘Pd J
souaseid oy} SMOyYs 04841010
-aid moyes YW (Ny 9)4H
ym =—uounjos ay, ya}

pun Layo UL anpisaL ay}
dn ayn ‘ssaufiup 02 U0}
-—njos Uumnolg ay} azoLodnaa
fysnoynnd “ONH .2U09
UL ANINSIAL IY) BAZISSIp PUD
apiube * Fy fo wes p ur fiup
Oe | ‘pyxppdrwa1d aq
mr YOCHN)PA 202707
“sha wayn QOH “ous yyun

ayounjoy 1 "CHN) Pd
‘@apritoyyo suluIeIp
“orpeyed = sureyuoo = SI,

“UL
jo oouasaid oy}
SMOYS O4NTIpP Way

‘OPHN AQ poyeqtdio

-a1d jou Sst YoryM
‘UOIyN[OS palOjoOd
-9s0l Y ‘uajzom U2
ponpoid ay} ajossip
pun 79 fo unas D0
ur ywoay PUD 1Q0N
Yn uorLod p0UWs
DXUY *IOTOD MOT
“A 8 UTA OSH
peysiyas.u0. §=4oy

UI aTqnjos are asatyy,

"HOOOIT UABLOPY = “IO*CAN)UY ‘sTVrssyyO

‘Jo poroy[y ore

qorys ‘syeysAio Mojad Ul aprsopyo oarndandosopyo

‘nay SMOYS UOT}
~B10JOO yojorA deep VW
‘SUIULIGM puB “O%Q%B NT

"LI jo
soueseid oy}
SMOYS 9384I

914411 & “HO'HN |-drosid yoryq
Burppe Aq poly; Vy 1°HN
-u00 SI }J "Ng JO ooMIa | YIM aznprd29
-soid ay} savoys o}B}Id10 |-aud pun .sayom
-old §=YstuMorIq-aoaA | ur jonposd
V ‘10H YR paryfipron |ay2 aaossiq
uaag spy BY) JoOYoRD | ‘aurozya fo
U2 8LOd DA IY} 3013994 PUD | JUALINI D Ud
‘240)94L D wWoLf 4I7%P ‘2;) |\pDay pUdIQDN
YIU JOD U7 UL aRDANgOS | YN PN xp

‘COME ‘°(HO){D

‘NOLLATOG *AAGISaYY

"anpisas 490]9-Ysin7q ay) Jo sayy
PUD L}DN UL Joa ay? ayossyy *a7qua
“nid Laups 'D UL *ONM puv HOM ynn

"d
jo osouesord
a9q} sMoys
suoIpoyey
-00 Jv—ndol
JO SuUI4SIS
-u0d 94831d10

genf pup aprubr ‘ssauhsp 02 aynsodvay | -o1d Mood VY
"101d°("HN)
P100u PHN) “1O11*((HN))] ‘NOILATOg|‘aLVvara1oaNg

‘NY jo
gouosoid
oy} SMoys
pe SE
“seg jo
aj dind,,

jo wor}eur

516 REACTIONS OF SOME OF THE RARER METALS

at an iridium anode in sulfuric acid. According to Palmaer, the red
color is due to the formation of an iridium compound having a higher _
valence than four.

REACTIONS IN THE DRY WAY

On being fused with soda, the gray, brittle metal can be obtained

by the action of the upper reducing flame. It is insoluble in aqua — 3
regia.

Separation of the Platinum Metals

The separation of the platinum metals is one of the most difficult

tasks met with in analytical chemistry. If the metals are already in
solution, the table on p. 514 can be used to advantage. If, however,
the metals are present in a more compact form, it is quite difficult to
bring them into solution.

In the latter case treat the metal in as finely divided condition as possible
(filings, etc.) with aqua regia; this serves to dissolve the greater part of the

platinum and palladium, as well as small amounts of rhodium and iridium.

Dry the residue (osmium, ruthenium, rhodium, iridium, and small amounts
of platinum and palladium), place it in a porcelain crucible and fuse for some
time with ten times as much zinc (or lead) in a current of illuminating-gas.*
In this way the platinum metals are alloyed with the zinc. Allow the mass
to cool in the stream of illuminating gas, and treat it with hydrochloric acid
to dissolve out the zinc; this leaves the platinum metals behind in a finely
divided state. Filter them off from the acid solution, dry and introduce them
into a porcelain boat; place the boat in a tube made of difficultly fusible glass,
and heat to dark redness in a stream of oxygen.

The greater part of the osmium escapes as osmium tetroxide; it is absorbed
by the caustic soda solution and tested according to the table. Mix the
residue intimately with sodium chloride and heat in a stream of moist chlorine
gas. Dissolve the mass in water and examine it according to the table.

_ If lead were used instead of zinc in the above procedure, the alloy should
be treated with dilute nitric acid, which dissolves the Icad and the greater
part of the palladium. Precipitate the lead with the calculated amount of
sulfuric acid, and test the filtrate for palladium by transforming it into palla-
dosamine chloride, and then into palladium cyanide. Treat the residue from
the nitric acid treatment in the same way as when zinc is used.f

* The operation can be very conveniently performed in a common clay pipe.
The gas is conducted through the stem and ignited at the bowl. In this way the
gas continually streams through the molten alloy, keeps it well stirred, and thereby
yields a uniform alloy. !

+ For more detailed directions for separating the platinum metals, consult
the work of SarnrTE-CLArRE-DEVILLE, Drsray, and Stas: “ Procés verbaux du
comité internat. des poids et mesures,’”’ 1877-1878 and 1879.

A
PAGE
NM ON ce tide ve 4a Sloe aieaig << sie elspa wt Rove evento? 122
Acetate ion, effect on ionization of acetic acid.................. 0. eee eee 46
IE MS cg 2t Ree Ih Sens alc Fla mow a's pig. «a gins Sut Sea's 00g ao are's a 341
I ek er 2s e's aia gia 6 Sa sist k costae oe ae ti elSic ove eg ka ead eos aes 341
jonization in presence of acetates................0-2 eet e eee eces 46
RIE a5 8 ory 'ar'a ore aarula aks WE o sk ad Rita ad male a talons <
IE eA EA ST oe a in EK Wan aah 2% aye ee sun wid ae aoe 80, 342
Acid properties, cause of.................. Soy So ale O70 AO Wea ears Lc pee ede 5 SST 2
RU OT ew eA) Diani gare hee big ga" aie 'b4'-dlemp 54, 447, 455
ewe, NINE NNSEMIGD chao cis 0. a gies ae pls al Wale © al digs ss 24. doug si eR prs e abe Sie 7
SSO RIS SPSS: pan aS a CAD rae te Pot NOR OO ae AE A a 57
SO SARA SASS ee eA RT aR Ra ig 284
OEE, ECR PES AE IS Ihre JT SR Ur Ee SUT EA OE Pe 56
examination for....... Sh UU. TS ROW ad tara aa eee te gE 284, 446
hydrogen of........ BA te is tat ergs ey emetic Wesley ROSE Seo. hee 32 Wo aA eS 2
REM Fer Eis oe deer oa ee Se ee Sa 8, 10
NE GUS Se SS oS BAT | Se PR king age AaB CORRES PoP tne Sy eR PTE 6
preparation of solution for analysis of... ............ es cece cece e eee 447
URGING CONCORD UTOLIONS 2% 5 oi6 os oars 8 oS hips Seaweed a5 Pare Saye NEADS ak 71, 72
MMT aRER VCORUCCEION, OF 65-6 6.5 oslo x sos 0 Llc deld whe a die cae theleeg bb as mie 54
EE eat ER RN ie ar Re See Gea eS PERE 3
ES ea eer reer ORE es Seen eae eter eae 28
I SEESIIS RAD TR Sr ECS Se an Pa a RP or Poa 19
AT SG yc a aici iu, wici'd kak Sete w cal Ca urs Se o's v oR Re veld @aralwale ees 413
Tl a a ieee ere SG eaRN SALARIES wh Saaz Ditsy 2° haie toed Were ae 82
MeNIRETITLE OF ALU EL yo oss Seis kei Helse fo ainke eursa ccd yes 0k delela leeis 105
Re soy co cidib a viv. 8 dal'ary, 6 5 Wea sce vad & io ew Oaiey Wowblgws EEN 77, 456
MERI atric a eR NC ees sales: Mas 86 ee aoa Eee oias wees 97, 459
III RORE TONNE O02 6 oa «f'n. v OG 4 vce bi sv RUSS Es tosvacies Wadee Reb 72
MINE BLLICHL URS J 524 Sin Was sds we 8S dy bare ae S Cao's oN gk een 419, 460
EN G3 Vie Sik Se PAW GON Ca dhe ee OT Ou sls ode a ck hte MOTE) Kee 77, 456
EEN SPO INMENOSIUM 5 6X occ cs ek loess we he Ho Renae ewes enaels 96
ROOD REN cn Bad SA ICA gik V2 xb didie se hig eed + Rea Rea eee 188
Rei a gin gig 0. ct cin kc See e ew de ae ach Wied Beale dade epbes 101
IN EI orn S wie 5k speck e WAGES clei vied dle in Xu AS aa ONS 111
Goparation tromy alkalies. i 2.68. i co iels cece suey 111, 113
MUON ok pissa oa CR aM Ais ee OS Ke 188
RTM SS ne Clara avid pa ask bos 6.6 ora HAS A eee kaw arpa we 349
REE CEOS Ws ose oi oe Fee Nn as cen eeee 391
PONCE NEA Oe ORG cas SOT TCRSY 90> dew oe Saber on wi 48, 447, 452

518 INDEX

Alkaline solutions, notes concerning.......... Raber x ie Siege Sm .o.- 488, 452 |
PTIATIU G6 4'og wih oon 8 ok PRO aw Vila tice heal, Cea Rea aie Bia era rena a .. 468
Allotropic ra0dificd tions. «0:85.54 swan cae dapidin. ald Mole Sale Wh a eee Selene 196, 339.
Allo Ve ARAL VGIS: OL. 6: oe 5. < oield.n sip hus Biinas Kiateias pl-nuaerac tee aaa ke he
Alm, 5. 5 55s won Soe eb 2 plvonlb yie/h 4 le allege eR wee Wamls, 5 eee WSS ak
detection in the presence of organic substances..................
oxide, method of dissolving «0.0/0. 5.6.0.0. Shene > 6s alee oe
hydroxide, ionization Of. .... 05.25. ..s4.« tia eS b> 3's oe eee ee z
separation from other members of Group III............... 189, 443
PEAR eo, CSS a vig poate Sle crass wis gies & salen Sy oe aS 25
PATGUNGBE, sisi soles as aves bong 9 6 6ieiwie,c BoM Aw # ol ebe oh © oats) ele-p aimne Stain a
Ammonia. See Ammonium.
equilibrium with water. 6.6.00. sso ves eee sos ds akce et en .
Mi rinking WAGE. oO Sis sie ews Masvimns Oe Adds, oo we 2 ee
Ammonium carbamate................05.eee eee PS aS peas as 6 We 3
GRED UDB. Sock waste ek ae we oh LN OR posed ee a Pe 72, 103 —
chloride, use of......... EPUR ERDAS NAAR RES Do 19, 94 —
hydroxide, ionization of. oo. 5 5.5 ou. Woke etn ees pe 10, 18, 46 |
ions, effect on ionization of ammonium hydroxide......... 5 ane 46
molybdate -reawent 605 Fae ie Ve tl ee 72, 231, 379
reactions of... .. GAs acs als tira ara bye as wattage ale kit aha eS . iheegene st
salts, action of strong bases on. ......). 6. 6) eee ee . 87 q
effect on precipitation of magnesium.................. vo. 94S
BERTUT GRO OF 6 oe, sy ats ole Bee Slew es ate eal 5 we eel ao ee . 92,98 —
ealide Brouiie si EA #4 as piaide.b\h e o4.4 a babe aoe een oe 461
> BVI OE. 8555 enh canes pe se 189, 443 _
Amphibole group of minerals...) 2...) eo nid citamiele eon oe lee 8 ago
Amphoteric electrolytes........0....ce eee cseewecees ate ea eu oh et nn ,
POR NNT GS 5 5 a. 5 cin soi hwo Wig coded nln’ p's ere thie 675 050g WAL Lemus teRIS oc
ial yates; ‘Chemical: 3.5). :3.5...5 sks cccaly sw ere be ce ve iv ols aD aleve aly Sate cheer
efuialata tive Sc 506s NAR ST yolk cw ers, 5b eae ee
CPUAMTIEALEVE bosses Sie este ATE. SEN Win aln k ie Slot a See
SPOCUTORCOTIIC 5 255 eos Bis oe ee he's's 0,0 pla bie lgps) «waa late’ Bee pipet en
HB DOMIB TAG. iw oe ios SS ewe 858 BS ale wee wis me ore A VND pha nn Wie ee
tt eee Pa ee Rac aeRO Dee
Anglesite........ AGES: SCUBA OE ESP aR 102, 205
Pat a nS Ct
SIND sis ais ES oes we ele ae oe ed ety Ws Se Ce ORs Se Port.
examination. for. 220i. cb5 RRS OPE I Pe oa
prelimimary examination for: .... 0.60 525.0005 LALO oe
renotions Of 60/50) ev eee SIS ANG Fae esto ae 4
Mmaboraites 25k ieee ab eas ORES eM ap ae
UIE. 25 02S SREB OE UNA She OA ov lee WS ola 102
ABUMODIc COMpOUNGS s 1.15.66 i i SWB ES WO i et
Antamonous compounds 26065065 0 i Le OCA a 24
PR SATAIOIAY 6 aos 0 5a 88 OREN 8 TAs OT RIE OMIA eben OA respi
separation from other metals of H»S Group............. \, 272, ann Se
ot) | re ee Pew wL rewire ewe ere ey
NEARED? 5555-50 Dood Boba PERE pe: 5. od oa ee OS Aa ey ee eae 102, 205, 2h
UR TORIB 6 oicecic candies es Claeys coedaunceeiPieted bv cree e epee #9

INDEX 519
PAGE
MEO MRD ee epee OVI ee eC L Coy eee ETN ETE N OE TE RU Ca we beige bre 102
Ss 2 GER OS bs A RE aS a ak ine, PR gD La nena 278
UNDP SUEY REIT A. CRS Si vs lores evo d Rye ew kage eeeud ase erp wee 4, 5, 7
RE REDD eer LEAL 6 yo2 ip FOR RPO ers: c-u/ata eh F oha dodo vieiira sd bie ge epbinleee water’ 223
ONE AE De Oe ep hard at nn oR Ee Dey RE ae era 229
RMR LORE TOE (FP hc OC ia here a ene 5 cated) va ee noe Obevis Ohne 228
detection in urine, blood, milk, beer, etc. .............00 cece eee ees 238
GREE ADOT vic i ciesccadesy buna oes Re Dae Ne Buy eae & 236, 240
NS Pe DD Ee or ar ge ea a fe SAE PE >. 1288
RIES PVs cr Re DO IORI a OOo. Ls Ape ee se Pe ewe 232
EN ESS evn ie eamren ly ek aia e Cats g0 ba Dae eb a eRe 229
Reinsch test for......... (1 ats, OEM mild Tatice RAUL eR ee Merges Oe ae ore ESD 240
NPERIPITRLL re Vic tic tat roil oo ol arco Ci ariph hint oh ha w/e Uo wc ave See av ed ON IPT RINE el 224
IIINPIREOHAT), GNAUV SIS OF... 50/6553 woe oo arn shelled Selle tie een op Sorel biie gk GT es 275
NTE BER IIGLONLI gs 05'p2 +51 os nese sist ap ekg eps deal al UU Tp ieee lu hel dod ete SOON! ghee 224
EET St be ls biter ok ee al « Cc Wn A Why oA wo RAB Re Oe Wee 223
EN bse S Si rie TGs tes oa ate MAMIE aie Sipe whe Wm MT wie sw RASS ered wale eet 224
Arse (arseniuretted hydrogen)... 02.52... cece eee eee Oat veces 232
TE Nero ey ea eae en 7h Seat, jelabe neh oe oe fe cn poeta severe sere GS Mo 93
NE aN 56 Fo i er vg wg Saga vc tote HED, thw dy wien OME Re Mee uttbe ne se cere le cr taee te 214
IIE nated Plies Sac dale ST sg Gk ena te lyin tm is Ho, candela 4 ve whales “Mle Ss 93

Autenrieth and Windaus, separation of sulfurous, thiosulfuric and hydrosul-
. IEE yo Ss eee i Savy hae tie Se eae Se Dis Wael. alae cig win Oe See 390
UMMM 0 28 che aac oes uss ss dnese oar tna pNNS ew ass Se ee ey Ke ya ee ey . 170
TT a. ah es 21 sae -g oy a dna. sahara ond cota ba te tn vou? lle) 0% ou. eR Sape ed Aes Wha s REN 214

B
EM SSO Coal ie oe atid og tweeters a Pele ete AGS eat wy a Fd SUL dee ee 462
IR Na hed. Scie oe eet ns SRM iol eh ETN inl as ee + Ge EN bo Vee 102, 108
EUR ST ees, AAS Fy 8a oe hal) fh GU 2 OHH a goa bon > RIERA A a UAE eee 108
Ea ha a ETE ie Sip Rta AALS By SQ CRE Eee Sy Rie ER Map ays ese 2 8 450
separation from calcium and strontium. ...................000% 110, 113
ETE RLOMSOOTPTIM MMU LOTS ONE oo), ise! 6 ua san loch Su. @20! olere) pala phe 109, 437
Bases. See alkalies and cations.

ERS ARREARS (ae, RAE on ete a Re lig ean 8 REL ia e 72
RENE SOR SRA TE MMe 2 ARP SESS Re Zee we IAL RONMNE BD Se o-E De ego 54
ER NONE st C1 CRA Ne ANT Pea Fe APM RR Se SMe es 2
a at holy Gey ped clte es Web Ra ahi go I etal 10
IR NE 2s) Saas gocts Ny aa the ah DEM Cera een do eueaw RE Sane ee 6
RR YE a Ly bpco edie f lfteterh ceed tor Sawlett 2
UND SS oy a pie led p bbb wiv bid cle op HUlpL FNS DUN wrk 125
ESAS RICE er CaN are a re 64, 425
near seem RT: GEIOTIIG: TUS, Sg ng vc vp oy oot sio'y t's pee ab pap a veeade eee 238
Ie a Te a et che aah ete eae avinte Ge wens 310
Ne NTE nN ee es aaleer ak Pa-aue 6 ba blarted meee 461
Ny PR Re he Nh Fe Oe aan wlohe a ide veh Coe rey ee re 8 461
ICOM TERIINE ER ol go) hE eg ule iy pv ua s 485, 496
Oe ee ey ee ON eer OE i ro 232
PeGhUcritetn ee MIEN TORRES re ok he. ay ecg y gene bee otae 228

520 INDEX
Pisnuth, peactiona. +6266 asc choc seh owe on entse Ob eure a we 1 Se
separation from other metals of H.S Group................... 272, 4 41
pabastrale. vias ak isa ae STE es +6 chen Said gales eee pe ad
Bismuthinites 2 .:..50.5 Fs eva bh daw olin we boule Shean ee a Wee ge ee ee
Bismitithalc ACI)... sso abs deeds ese e aes ya
Bismuth-sodium thiosulfate reaction................ cece sees eecescecees
BORDA as 5 5s os ital we as so ioe als eee iy Da ae were
Bivalent acids, ionization Of; oo... 6 ses a Vek a es ca sem os eae
Blochmann, strength of reagents... <-.. 0.6. 6 NG aol oc =
Blondlot-Dusart test for phosphorus... 2... 20:3... .s0s ees se e003 ee eae oie
Blood, detection of arseiiie in... 0.60 bee de ey ane ds dene eee ee -.
Blowpipe reactions. 6. i s.0 a ee oh en a Paw eee es Oe 68, 426 .
PR eo 5a 8s a ae lo ate ee eH pew thle ON io ink Pee a sepa o> 25
PROEACID DOIG. Co 'h6 oS 5e s ae ee hw ko She Ree bee ee eae ev ade ce ence
Borates, solubility Of 5.054. 62 Sia 6 Saas Sats sen bo ptace ce iss a ek we
PROMI Pe G8 oh goo Taco k 0 CA le a eh OEE, Meas ee 82, 357
BMAD on, oss cages goatee Me wesw kerk baie Bes ere eee ae ea 64, 425
SE BOUL. oy ac 5 o koe eee ae Uk) A PhO Oe aE ee ee ae oo sree yin
WPRRIGG 2 Ne a ns Bg oS Daa ee ele gw die + © apd aaraly ileal dee fia ae ann
Breakage.of dishes, cause of, . 0s) vs ov. oes o's de oe 08 4.0 bn oa gk
PRTORIUITE. 5 50 FSA Re Seok aa os 8 we ah nes pb seary 9b Dt ae
Bromine: (free); reactions, . i666). Ss a os oo tote Saws + a > 9a
detection in non-electrolytes. 0.00.06. 0c ccc ceseencesocceses Sen
presence of Cl and Tow... sc ces es late By ceed Se
BOTOORTOBE So Sa 2 05,3 SENS coca 26-8 Male 8 lp rin ada Rela, LAPP Os eta
BRCUUE TORUGIONL: 5 5.5. isis) o<cheoeccBe ppv! atnle are cop FE AE koe oer 333, 395
oo ost) i a Onc eee eel MME DEINE ER ara gtr ten
SOIOWIES ids ob bo 5s agi TE aie bw Swe Obs 4 eae ae Cee PR ae og ale
WATE, ches ois SSRN SS Ma a uk Tetwike bw Woe owe Re a Oe ete 116, 284
Mame, parts-Ols J. 5 sie eo on FG mw bb ob eee else dee 8 Da
POMBE. SS ei ie eos oO aa ee Pie a a ee wk we vie Gee eC Cen el ee
C
Cacodyl oxide......... END kG ine bs we a Ge SD bs Sen eee <i
CML TORCHONE. ook cools 56 b9:a ee ead ee oss bb es oe
separation from other metals of H2S group.................. 272, 441
ammonium sulfide group............cccesecveecs :
Omamim, Teactions?.. 0.5. ova c' os bx srs.op0 > vive bio biped alg a 2 be dl ee
separation from rubidium and lithium...................00e0eeeee
MSNA 66a a5 5 a0 Sigla pine le 6 4 U oo pro Noela,ale idle, ¥ Rye Sieh os Sakae te eed
BT RR Se Pe oh Oe Se Stee
IN ob gown 5. bhi e ten LAM acato- te shih 0: bcos oilaghe pew LS 1k Sa ahh lg eRe ao
RBI, BOACHIONS. | ..3 5 bew.ow So oy e's 0 seo Veoy aya rae bees eee a ee Tunas
DATISNAR Ap oC cicecabed thin bin! s, sracach coe bes Loa habeeca eh kee atte eae ge 105,
WGP oo oe FS GR iw inn oin eo Jie one wip nas Oo tele ee ae
phosphate, solution i0:ACId8.%,.. .0 + is ecce ceed e vy oe vic nls 0) SHR
PROSDEMIE. 23s idk adn aaa biganseehe¥s eek se el bie 106, 439
separation from Ba and Sri... secs .c wis ss hb 0c Ue suiels oe sae
ROMEOION ois 4:4 aishe v0) Holmen besten. ovo wlaia bese Oo tint © DR eae ees Me oe

INDEX 521
PAGE
MRI MIWM: ORUGION. OFF Uvshs sc ccs oe tc ccc cee ctcs seeders eens Sit sa oae 439
Rane MTEID. COLOCUION OF o..04.5.)0 cole se eicieleeca's cc tccueedreewieecicessed 220
SrERONLeE PCHAVION ON HMNIUION.. .°.°.%'. c/s a's bin elec dive peelaneceeecetserecce 355
Palapiity Of oS ie ee a BP? Dare N IN Ere ie te 353
DerMnnNE Men SORITAGIOT OL... 2. ss wa 2s s eosin ep ooo ss cbeeeetteneses 8, 10
Oe = Ie ae as Sea aon mip ahr) Meret aD ONE Yq REET Si 352
Sere Sc ee Cea gS ein SV CE Cw tol thee ew eiles oe 421, 436
SE ga rar or gee ea Re eae? Ree 78, 93
Gimpociation Of... week akties PML EO eeu ce os oats 16
Semen OeeCIION. OF POLASSIUMD. . .. 6. eel ak ee ee ete eels 80
III Pa or Poetic ie oh s Peas eR tho d's Le hades Bless eta be 487
REE. oO. eee Ss Pek PS Soot alec ghee anes 403
ET ges eas al Suc wSlag ns Dlg oe hehe tg oi alalg so 249, 437, 462
Cassius, purple of........... PARRY st CLAS ETS ROO N at ces es oes 263, 265
MOMMIES ee oes a ce occa es cdot ae eRe Re Tea ee city eWLee eee ae 5, 10
ERs rs a SOR I Tea Site td Ce ogee sees ea eeudeas 5, 10
SINE Soy sos au Die 0 cals soe noe wis Wea ¥idlee Si Na Win tg Vinars tals aie 77, 440
Caustic alkali, detection in presence of carbonates..................002 e000 ' 455
EE Rt PO ee ao 18 Se UNO R eee de SU a oe sels bee es 102, 106
RemrIeONTIOUTCS “FORCUIONS 62... 21k 5s ek ek a cate ete ee eb e eet eenene cae 470
MEMES ry. cla ele Gas ocd, satee oGjd kde 515 wend Sere eb Ole wy a ee 478
SRR Se S Te Noel ees UN a GA's a ek ARS oe ate 5 be adie Uae 469
re PE SEN ea FY Sia eel coc wielelaierete’o esi utelaws «t'ee oe cea 469
Suman TOONTICtINCls, PEROUIONS) 22 62.2 kd oe Sb oe dice Loa ge he ede e se ots eee we 469
RE SIE Los Siw! ol el hw Sal vhs Sto Kw Wf dela o a nvoralele eyelet ale 102, 205
rr AS a Pe x aie Sa oe Su scone wah soc pix ahd. toh oc SEE onln eRe w vised 413
EERE ag SEU Scone arn a Cae ee RE Py ones? BEEN S|
RE eT ek hs Dar inthe Dal alata wi alate alg’ gtaitiateriigj Pod Pawan 0 ER 205
Charcoal reductions before the blowpipe..................502 eee eee eee 68, 426
IE PERRIN ot fF 7,4, Esk o vke eB ose uavs Uwe ly wl Nee. alee A ated «6 65, 435
tar iar 28 ee AEE ORS eae med we ee ead cae cehede es weae s 214
Chalk. .... RT 1 AO eyed hse) Lube Taw acs Ace Mette eaten 102
TS EIR 7 en ora 2 SOP ay I PRM a ert eae 1
IMEEM TOERNA S os) ou tne tac cig’ v6 a Sick Ton’ vid Oe caida ata nod d'é.g'ww ge arnree 13
eres et ed Gillen Dy Sk ev ag bt a Dees eld eae aie 3 Ord wiataieie «ere 1
5 LU. idk Ay ck asd atthe a ae ands ofp lu's v0 wb ev oe Pate e aelw nee 82
I ras Scie gin Neco gis a ahh agave este Se aaie o cidlaleaig’e slaidieraty dace Mad 172
8A rg rahg. sgl an grePanat baie oes 4 GE NIV Wide as Sate Bele Ba 397
ERMINE oh oy hho ois. ab oka woes eials’ AelaaL ew. cha’ nl ed ¥)0 ¥ PeeWee eo A Ble werd 397
detection in the presence of HNO; and HCl................... 398

Chioride @etection in thiocyanates. 2.66... cek eee eee ccs nee ece eee ea tue 326 .
RJOMOTIIGE EHAVIO’ OF UNION ic. bs. etd ceeds cee ee cc eet saws eta paee shel 290
EM SRL K LG) 5200 5% 05.) Sed of heed Seach y! st ah shainte a shin Gy voala: the Whe Wig a 5 287
ea INET MEMOS Sco Sg 4a aie 4sar so ie when wis a sleeve voy Beedle ww hes 291
SALEEUIOM IN NON-ClECEPOlY LEB. ios Sod cree ev oie ee a ee ce meee eee leas 290
TPOMOUIOR ETRE ENCE 6 ose oa ho se 6 sles OR eager a wis ee we om 306
NARI Ane arity sletdiale Slo & S's eS lek us See Spe «OE 416
CHloropl Atte AGH (GORENG ele o eA ies hae bee ve see rete seed epee 73, 269
AAP ORUUM RUIN Ae talcta ps nee vies Stores os Se PW IM ate ne 6 0 Vole Fv We ONS 8h ee OEE MORN OS 135

§22 INDEX | :
Chromic acid, oxidation by means of... . 21... .. eee eee eee ce tere teees 32, 138
eompounds; reactions: . 2.0)... 6 Rk Se ese eae +o
Diroauts ee eer eae PEP re te sae ANRC . 182
rig lyeri tk SF ie vicky Sex keel Odeo acs bmiteia wahleaeee Ohad pe 4
NURI in 5 Pas Sa kd odin wi Ca viewed ep RR TENG cain ga 32
X20) <0: (ee 8 I ee ote aie eae ag aie ae she eee 182, 4 a
separation of other members of Group III.............. 189, 192,
Chromophor er Perr it ORE Ser IAN re ES
Chromous compounds, reactionsS...........:00ceceeecuceccceussceeunee 132, 4 436
SPAT yeRIOR GL oy as ela Sinieig cede c's od cco deo 6 ob pale Uleuible o cneeee tie ean :
de En ogee eR POMPE AM Gs bei yE eC Le ori. (195
Ortrates, solubility of: 55.0618. ee SCE asec kbc duels va CON ee
ATION occ Six, Sas oa i ie Cie pels, son 3. ca awe os oe oe |
BO | ee Sms SM Cee Mea OSE TAU A TGRE MEMES SM
RABMGCANIGG ooo. scinropacy, eoleie. boas hc ¢ ohcalond wodie gro. dsm Wongacn ile maiko Oba uae
A es a eh Nav Wb res cresan ad hemes ARK Ua esl Co ak Day ee oe
Closed-tube test......... Hei ee eOR oR hed Bk Die Nee URN Rane oan ea 64, 67, 423
Cobalt,:complex ions of... 2.4... i sees ie eee eee Nya, Se ole eee see eae :
Goeobtion iy nickel alts. (525 305 cn cienic = ohn + aise hes Aoloas - 0.3 9! 0/, 184 at
hydroxide, acid properties of .............20cceeeeeeeeeees oes 180
i. itt) A eee ee SOMME a Um SIGKIIS 178 .
separation from other members of Group III............... 189, 192, 443 —
SPMRIOARRCTUMANI GPU... 5 sc ou bat owen Roos Moe hao Pee eR NTT sate 181, 328
AON GILG G) 0.5 aie soo; <ieco.es Wp ee Ne PN PR TA ON EA) eR Gs) be se ajalais oases a 17S
Codein test for selenitm .... 2.3.6. odsi eed ecsscs sues seclawoe cena anne 499 7 :
MONA aL IAL IOTAE 85. 525, C niyo 013 sic yes oid a bele's piv elec ere geolelend eto ea 58 ,
POOCMDCRH 0 5 gk CAs Roaleia a a ee aes 58, 127, 218, 415, 495
MEMMRMAA Co rsits: $< ooo: Soh np lw wn eb ol opp mane bm; Sup hanl Hes REA alk tatoos pe 59 =
oe i a eT ATOR eat mre Se 60
Color imparted to the flame......... 0.6... ++. eee e eee eee eee ee eh y tas 64, 427 —
MAMMIUAUAONO 5555. aio gb ve acc v6.0 's jaaete op tis wip» Petia. 9 oss eld Siena So ela oo ee 499 |
TNMMNEM RCC we vulid sv ie: co poo, vinta Voir Wiwtohe soem Sb (kot Me vo RST ete 480 —
OMT FON OFS Cb os sok DOSS os ied ws wera tw Bhp eo ware atl OO 45
COMATIOR TOTS. 25 ode kcoicicie 8 dip & ods eerie ¢ 0 di8'e.p wae Vale epee ea ts emus Sn 24
MEMNSOE OR: SN rsia Jay fin eee re oan 254
ealta'and double salts. 020: 6046.5 60 soe Bed soe 20s Oey oe 25 |
RIMMMCRLCEAUION 56655 fv oko ovis bolas vg so dies nlp ocd Had ov a a phip wale ile 9.
changes and ionization............. aca-0 0% blk. pele eo ae 18 _—
mole i Sy Soa deren vie bin 8-9 8 0 wie Shia eb Wel ghar 9
TMOTTENA o5 5 so na'nce+ pele 4 © einipee eo 9/ore lp binlleee we elegant 9 =
Of TORMHUEE, 5. i ies) c.0's + a nwe.n-o Gi8Se win} we ad ee Pe
Conductivity and ion concentration. .. 252. 660. essen becky egestas enya 83
Conductors of electricity... cis. .i6000'. sy 6 4cs pele oes 00's s 408 trie 3°
Constant boiling acids, 5 5 6\)be5sc's4-0 + + Jovot ove ¥e be ¥ 0 Unicel 57
NS S57.) bs co aie ac Fk Vd aR eh pws Soe Le eae io 46 2 ae
separation from other metals of H»S group...............+. 272, 273.441
(NEL OHS Group 5... 6-03fa'os a. eater Fe ck ae weve BIZ
A RMINNBL Sooo ig v5 Konus! -9 4's igholoverWiaCdre's 6 4,ug ala kesh op ean dialy oh en 125, 438
MPN RU ORCA 8 as. od oy on oe hie Up ive Sf eve ple toe ea a ews gtbr ie ol ona 362
MNO EOFs 0d slo iad Wea the CEs Dame Sev uae Phas Pia ee ree e ee .. 132, 205 —

PAGE

Crookesite .....6..+.006- ES ES OPER SCOR Cha Id Sareea th a a gee RPO LID ee 485
AMAR irre Tan gd Sethe Sin cals «vos e-u's atk «5'9 vaio KKs 89's PA ie ge rare PEEP 82, 125
NE ro aly seid ies sca se oy bn ands Ves BPNRe eee ee cers 152, 160
NIrILOUNTTONITICLE, TODOUODS 6 62s. 6) sacs esc es teed sa eh c6 aie su bend die wore 217
ERs e te Cn ta aN so ootny Ag ACRE Os DEAE Sek Re Siem bleee Heke 220
eR bie age bik. Nak bie cao le Veu usc ees ka cede Rae Aes ae oe 214
NM gata ici oy asp. shee yy 4 dies 'e be fv d.92a'9's bine Lime CR ATER N fe guna 494
PMI OOMIUIOUNGS, TOACIONS. ,.. 0/0 cose cc dace tess tedvecceevccedsvdues ear 215
DRE ste wna a psn Sie sees « Oo Pei ohh eid Gide lbmrdiel gi xa. e ce cree 220
ENERGON Sy ee ek wd Slyia wae Eh hs 0 8 tp ow ieus Fads hv aes Sree woe 342
I eee en Sino Wich ne ae «wile Oe 6 ghia ae Si Kia'seg Wale wale ee oid wD, He 342
EES a Baer, ee Ot eee Ae arn oe Wt ree Simmer a gas Gry 311
PERM NES SOME G UIE olla lc 5 1- ciag See aca Wiebe ode Dio a bd CO odin Oe yep 315

complex compounds, 24, 146, 181, 188, 219, 280, 312, 315, 318, 321, 328, 438
decomposition of. ............. ME ee L DA: kee CEA OE Uae 152, 322, 438
Cyanogen, complex compounds of, 24, 146, 181, 188, 219, 280, 312, 315, 318, 321,
Z 328, 438

compounds, decomposition of................4. 152, 315, 320, 322, 438

eT vo ca oo ae Cay HR Gata Ck AEA KU db ATE 6 Raa 42
a EMU EIUIUNTACUS TORE dhe cy 5 sh cicis vile bn on Gia aD ev eb o0ik him o Seed 418
MN MMIII og Os ake wy oy la oihckS hg 8 Ree Potwine mv es hoe ee 6 Pie ie 367
EMM A Sek. way eae oF whee. Spb rate WO DUNS ele bd hin se aeue 34
IE RTI tastier aoa. 5 la vie 8 ak dana ai wie CRG B-o-4.6 © lnceyn ish aee Bla gee deen 516
TIEN He ciel vata lc o's ahi cies woh ev aft © oe Maes ROK AA RTT ee wae yes 125, 142
PIAL. ACT ATE DOGIOD Of. oy ics neve chakecdbecev des atuwdetebee 32, 138
IMM ial cary iso he Gok We KS wip View a pono) boa, Sw Raiecln 6 SS te Se eat 316
IE Reg gra sslu fas ald. V6.3 alow 04/0. 04 & bio apa. wie Wey pine oars Con Ou aie bk 469, 473
Diphenylamine, reagent. ...........e0..c2eeee- Pte Eeyore eee 332, 394
RIMINI Ce 5d sre ae go Hic anole VL oe we ek SMe NWR ee CE Mee hoa oe 177
SIN PRO OR ULE: Soyo) 5 Uo di x bc os we Bae ass OC end ee Lo Va eke oot 4
influence of concentration UPON... . 2... beac eee aise bows 18

BUR NCWANTMCIINCS IUCN 597 os tivo tk S'Atl ap Sins eco Bo wk eid ningeoe’ « Mak edee bale La 8, 10

PRMEMIMETER Ar Sees ee en Neh eRe ee Wake ea oa wae end Cras Min ee 16

EOS ce eck cis Py ee ke eet s SP VE Soe Cpu bruno 6

PRCISNMDIY PRINELO A. 0. Se Wie lei goes Ces aie onda een eae 10, 145, 334

SPIOEDOLUE DONTAODIONIAG. 6 Sa ws vhs 5 0s CRN okie lees 16

MAANEE, CHRPE OAL O55 5653 9 0 Gk SiON woe Mala o kic ake & eee ra een 9

MES 6 oO cceini c's Sateen d dee Conk Poarhww sg Me Vis Kine Oe TELA OE Seas 10
TCS Si aie 05 cv tis Vale aU aves dbs od pea kedbaaean 17
NM Shh Sai sg VAs oS eld Male Wei oe Pee ee 17
gg Se ase sais) ie Alace aero 0a woe wd eg eae Coe Pea he See 171
PiPe eee IS INTO PTOUPS 6 6 ons hoe ca ee gs we bee eve wh ales ews awous 284
DN, IOC EPOITIN og CSc di ak pri at bo as eb @ Chal owe UN OD 70

Re fain y csc Vales oe dees aches ee eclene plat nerd chs Bip 93, 102
BEAM EE ASTRNCNIG BEL Pe ore one b edhe oc vba sabe h eoele d amwlemd net 90
MUR his CHEE ME gS SSK Or5. c Vg wes oss ws a lnahwe Src obs lhe wate te 332

HOIOOUTIO CHATS ON IONS, Fs dics ss 0's v.o'0.8 acactse Minions plasain-e a ad boy sials einen ahaa 7
Electricity, transport Of......5..ccceccccescees evoke Gaecb Glaus, Ws aoe eels a ae 5
BRCOP OG sina nies cog A owen Soochow et eh cm One Rea eee IR RE CO ea te
Electrolysis, explanation of. 2.056... oss ee ccs 0 eke 6 0 a0't 5 ce Pelee Sian
ADO UERS, 60.8. ccascla asked Wee eh hoy 6c hae Lice hee ae ee ee se

Hlectrolytic dissociation... ..00..3050. Dc. Shae e does sae eek ane ts enn

“Gnd ByGrolyels..: oss bc cceoes oe PP ‘ceed
Hlectrolytic olution pressure. .o5.5555 5 jose caie ace pacaccceb cap oceee ‘othe oa i
POCLEOINOIEOS BOTIOK. 352s Sebo bred ore cepa Sculls oon he ce ls ed 37, 41, 43

MENS OOUEE CS STV aod wey dk we kk Eee REEL E Se CNT 8 pag
Energy, free or available. {05.05 55 Say ds ase a eels oe os Pen
SE Ce PRN int Dab ra tin ake erry ere ate oe Be Pa, Fh ‘6203
BOREUITOD 50-5 Sic 23 Scr ag 8 40 e'eig & & ORK Se Cibie oie let ce Gn tuly yO escheat ann
Equilibrium between a solid and a liquid.......6..... 0. ccc cece cent eee eee

two liquids... 0.005 64.cs 5 os os eleee

WOWIGPONCGOUS. <a. Ss wile bw Gk ote wea bn b.ccdve wy De eh
EAPOTVRICTIOS 5 50 Sinn ohn a c's se Mek ae bas edd ebb. g des 5 vg Sah are Ol eee en
PPOIGM FOACHONG. 6.425 PSS a aie bee ewe ee ke ee a eee aye
CS 2 Fo Baers 0.4 4,6 aaa ORS ae We LM PRE E Tie ken evi deere aay
ibabstrse Rea ooo as PG noes hob bdo bo bao Pee eb ee eH."
MUGHYL BOCCRES, 6 o.oo Sin nce aie Bou woke Sa tea be baa eS hee pce ee es eee

maarise Citalide.. 6 3. ise Gs vs ores bins wee oho che ss O0Oy kee
Ethylene platinous chloride.................-. Nis» whe eS leleve, e hecy a
OSIM Ts sak os wee oy) 5 5 ii¥s50 ood Rac are wa able gelato 9e eis theo ae

TAVRDOTRUION OF ACIS 205.555 5 oon ale ped cei es cle» oe 6 0 'o'g mb sind 5 5 oil aie
Pismenination, prelimiwmerys 62.66 esi vhs ols ded beck es ede ce ov cae ae 61, 423

2 er ae Pee et eee ee eee PEPER EIN TROT ry
Panne 6 BOMON ss 5.0 ipso 5 6 o5 65 ak do o's saws bc tw We ae a ee
SERIE Ge ac ee Pee NERA Ree ee 78, 125
SORTING Fo 55 Ve Reh aD iin vee ee Fs y BOOS 904 Wee we

WOETICVANIC ACID. a; oa che Fake Ee bi inks eRe ie Ae ee eee 321, 327
tt a oo a a Te ers Were eee ya a
decomposition on ignition. ........ Ve ree 152; 322./ 3
EMSC ITS CN 55-0 d 525 tiene fdaiade so! 05 Siw ee 0 Vita ha ine aN ot sc pe 318, 8327
PRTORVODIAGS 5% 05.5844. ree GW gs 5908 5 oe bu Seok eae aed tea vende een Ss
decomposition on ignition... ..........0ceceeccceeeeeeees 152, 320

Perrone Compotnds igs .s oe ha cs cascade Meee rede Raed ices 6 bole 144
OXIA DE s..ios sieve sataas base Rea , 29, 30, 31, 32, 148, 147.

INDEX 525
PAGE
Ferrous oxide, detection in presence Of iron............. ccc eee e cece eee e eee 148
sulfide, theory of dissolving.......... ee Raglan AUER pert Wistnl 145
I ete Scr ele re no iss Wie-y ole.) Fe ee awe elk wee oe oa ES c Bee 58
Nee eee car ee, Vale aie e'elay acy v's oPec ole whe wo Hale eg big Oala'ge welt 58
Fischer, Emil, detection of hydrogen sulfide............ 0.0... cece cece ences 336
I et Fo aaa gee plein gents vine Ovo eke vewee ges ormie 64, 427
ETE S Sl cleft ear PaO TS aa crag pt ale oo tian ald o sree nis ea eae waa 62
RG ee I See aide ee porate unre ne 6 61, 427
NE ce eT haat, sdb er sas 0 Gor re RO ee hel dah 6 van Pein erga a aig alors 115
NN Fo 22 0y giv lara rats bitchy e's pc @ acacels egies 8 WS + eset vee, 2 6le 410, 437
ES oy SSE Omi eae Sire 8 ey ads veo date Weta ang om 407
Se aNNINES ST AMINECNTELON BTL SUNIGALOS S = nL s oe ce renee ee eb bee nen eececenepaaeee 409
PSS Grae ae ret nage 4' ate alo elated kak Be a a bate wie’ Kase 8 des 102, 406
WOTIMG MOM. verges ccs ee PR aeare ae idee ard etal Me a NIWAS glee so hess tata a og aOR 311
Fosterite...... Re eo ee Cog Mie as ih ne ee ober oes cxetnat saves eae 93
IID or PCMAG Sa a pial aon eS eee ba oly py aetec a eb wwe 184
Free alkali in presence of carbonates............... 0c eee cece ee eee ence ees 455
RR re Sans We kee fae Vis ORTON REM oe Dae wikis ws be vee 2OL
energy......... ES PRE RE RROD yee ie Sef ORO oS RC OL PE ERM ae IPE eve 140
NE Ee arte eh cet cig ig glaigtr cate nV Ok wiaW staimr ele aietsce cs wn gee sae 121, 122
Pe re Oke Cea Vico P RG hes) ebay OC Wedonk ee ee. 263
Re) Re ener gi Te gS clea Eg Gi oe dio SMe one We opeheis 4 Geom . 279
Rie ee ecm asie Cate Sirk sigh alsisle Sipe hae oon pb wie Sb oH 0 wwe 61, 63
G
I ate. eine Oy Net ate Gia d dese bowie y cutende wk Gea ws 465, 467, 478
ET Be ere tell peg als sy Mey Sarg Verwle. dre k hid id ee Tag emia 0 ee NY 205
eee ig cig Luo e pA WEN elder Aue wocig eS WU Ne easy h 8 ea UE OF Ute St 125
REA Sere ee eS ONS SLES eR ues Dene v4 Law RE ee wae ee 172
PeeAtION Ol WIUMINAGING. 6.6 Ss. a ees we oe views se ook ed oe ples 61
Re ins Sn ie F< SG eas wee Woe 'n ¥ eme woke da end em alet sales 122 -
EN ET tne Lao Seis aye ee a's aid ee EW A(e 0% 9 ROMER E OEE NS 310
NED 5. S007 0 oct C77 5 sas be a oie hee see yen ss os eH ang Sate ees 1
RED RTI ogi ton Saar gh wre ola eer ae we aise aoe WER ae teams 440
See oe SS Ae e's Sac Le hw Uaard Ae cw tale wi areed. im Cote eee ee 172
Ee CRs hg ond Sc Path inns VR aS ree en po we Re ese ol 461
Gold, detection of small amounts in alloys and ores.............00.020e00e- 265
EN es G0 Gh cS Sole 0 vie hast he Co CA Bogus a senasipe URS Ee ee 263
yt ca foie Liga 6 SHA aE A Seale eee ES OOD. ros wooed 261
ENC SHOTIMNIN 6 it Vie ook So eo ys els ie ca Cokie e he's 270
ahr 64 oy ae, alain OD oo nea Fw ba OETKER Hae SOE SN abel 142
TS Sonics Ee Le les e arssk'd oS ie Sa aks oe sag aa dR Meee 9
aR a BS aT. on Sd Ski's: 3's a np loug os Sv CM A Vee mee ae bie ee ee 221
Greiss, Peter, detection of nitrous acid... ... 2.2.6... 6. ccc cee eee eee ees 332
Rees OP BOHN TGO Co iice. 6a Seale ws a oe a SG bec wiele he a cites eb ae os 284
eae aE SRRMOE INES TINY G09 oy gk So ora 4 Gane. ose vie, co wUb Ie s Sie oY a Oe ee oes 70
Guldberg and Waage, ‘‘Law of Chemical Mass-action”...................5. 15
MRANC LN Mas TEIN noo cli eA e WN C0 wh. wb W's Aw 0'elb 00 Sep dw prostean Sees 238

VDOT bai Oe Ry FORD 000.09, He ees bike Viele aes 2 ee F.9. OO "Flee G) OF8,.070, OS 102

526 INDEX
H
PEE OK a ats ale g oe eK Basie aioamerg AS yee eee IN Are |
Halogen compounds, insoluble... ............0.ceceeeeeeeeeeeeees iakeas er
Halogens, detection of HCl, HBr, and HI in the presence of one another..... 306
in presenoe of eyanide::.. i icdalieniiiea kak dan aemcelecbene ene
ThiIGS RD ALS 5. hh ied eh oe Feet a nip lee eam .
oxidizing action Of. 9 6. 6bs bee sue ae wles «8h eck oe
TERUGUIORMUG. Oo Bee nnd Shes Nee ae, Hay eeee os OER aa Webi oe erie
UDA, Ee Sain Eee SN MSE CES rep
of formation of ‘water... 4.6 60 os 0 Sw Eee lp oo
ES NG So Sega See Renee Date SE Ly ae Rae Aces gists 6 souace sa 'e-5/ aot een Le ; a
EE ESR ie Ve RE SRR Se ees Rak RR aN te 351, 405, 427
Homogeneous equilibrium. . 0... 00. essence kc eee os eb nos bie own wie Seen
MORRIE WF ton aie da Riek Gey Liew Ho eae aisles Wa sack oo her
RR CMPNE RR ease kin al carn ca eR a tbw vole ce CNG ah oe Ute eres Ee pcs many Sos ‘oy
BRUM NICS a yet 5h so 50650 pu fw boy VERO RAE Tee ote oh Seng ew Gv
DAMMSIES aos eS vib ne iin jared Labo Btatioiee ous ably peel «aon ys coer SON Tealh WA S c oP
TIPU IND 5 6 oid 5 2.c ica we ssbe ase. chlale nex eine PS shh he at a I as
MARINI FUCA 55 inl sip ev has eas rs RRA Sve ov eto hv eae eae Che bee ve wines
detection in the presence of HCl and HBr.............. eens
POGUCTHON DY. a. 6a) Yessy wise ioe, 0d ek aca pak ad ale Rene an Oe
MESASPOUUORIING ACEC aos f'n eh a'y ond pare glee 00) Reader lott 6G a kh socal toe rede ! ‘s
detection in the presence of HCl and HI................. 306
Piydrochlorie acid. .~. ii Seis ob cn a be Ss ped Op ae oe 285
detection in the presence of HBr and HI................. ‘
HNO; and HClOs............. 3
oxidation. by nitric acid: #0030. scs's see dels 2 3p ne
POTOXICER: » 4)! 5.<va ey ks de sie ORE oa é
Hydrochlorplatinic acid. See Chlorplatinic acid. |
po iter ia fe nn: cei, aa RPE em mre eh :
In presence of halogen and complex cyanide.............--
’ Hydroferricyanic acid. See Ferricyanic acid. !
Hydroferrocyanic acid. See Ferrocyanic acid.
Ie SSM SEEALON ID MUBMAT 555 20. oo ch cp! 3-64.» 0: cg Soaiutin’s ehar'o apo. 9 Mee See eR bs ea eee
Hydrofluosilicie acid. See Fluosilicic acid. a
WIRE 2 oss ay 555 hsozs octes ov. a nk otc 0°) ae calls kb eee an 127 o@
Hydrogen ion, effects on hydrogen sulfide........ 0 ...... eee eeeeee 47, 145, 187
Oe ROU her ee Siok, ue PAIL ue eee baes + 6 pls eae 28
METOKIGE ee AS i sige Rais ache Bis nb ene ean ee ee oe ig
ORIDRTIOM DY 565k waic ORD Ea hails eos 3 ee ‘ae >.
reduction of permanganate... ...........0.cceeeeeees 85, 1714
reduction by. .5 5.4.0. cs0s caved sles tals om seein 09 el nn
sulfide........ i avapal ade ale ajak 4:0 ¥, oC hone aiee Se deeb bie "htace, een aaa 334
as reducing agent... .. i. kane sins bales 5 oe ny ee 35
detection in presence of H2SO; and H28,03..............-. 390 “oe
effect of hydrogen ions on... ..... 2.0.25 01esdeoeleee eee Fae {A
AOMIMALION OF Fis.) 5S os". Fa aes steele es a 10, 334
theory of precipitation by 0.0... )5 50 o.e feds eee 47, 186 a
BUR OIES oc ci05 5 ow SEK hig Oa ts, «oes nce aS HOt Mar HH nig a ee a, Se
and electrolytic dissociation.............:0eseeeeeee ve +0 60d 4.0 en

)
~ Se
& a

PAGE
Hydrolysis of ferric salts... 1.66... 0 sec eee cen ececeees Said ata lle dead ate ere os 148
EE RAINURIS RUM ob ois 20 a alld 5.2 i ldwhe cae ide! avid. w peerace leo p.dimdewverd «Wee o 158
PPPVENUON OF. sia se lek Alea beibie hia 6 Siake; ee RAN S o SLa'b9t 51
MTN Mm FE Min eet Rec 9K a)! gan ein °b 20,4 dala pga a ha abe a Baers wis esviehage x A 127
TNO Sc G20 AY foe ety cde Ua Wisin bac BU doa V obs en oe ov PHIL D wD ee 334
Hydroxyl ions, detection in presence of carbonates. ..............0ce cess eee 455
EE as We P00 wits ci 5-4 ig Aro Gig sluce hw bdo Sel OOS Ts ded « PO dwn c dhe 2
Hypochlorite im presence of chlorate. 0:0. 6... cece eee eee eeececccces 397
si aiy Linx ¢/os, esos dibs: (idhe bein Vo ha Side cares View denen ae. 294
PPV DOCIOLOUS BIG 66. oe ce ce eee Gamera g sas ¢ aha, stele heap AY tae 294
ENE eR IIIGY Of. Oo os Faas eens bccn cceedes cesuceceeteneas 345
NE TLENGD ol og ao yok ok" oi ale pcs sv GAleie 6 slaves Fe wie ae bee eae us 345
M86 WG va eiy sack wah ok kash cvdews sued cewen ep hink 487
I
a a ato 5s. 82 0 9 ela a Ale slo vibpepele, biace/Sa Wane Rat J wage’ 61
I ets rcs nicer ssc, p-crake esac ee eae ese aw oie alec sass le Myhan we'd cod 157
MeRRtRTMINS, TWHELOING ANG OXIGIC .-.). 0... iow co ee te hie cee elec encase cee 68, 427
ELEM OR ee Lo ony og Rectal dey Merle a wiselp Bia cme» weiele ages Oe 56
IRs re Sather s, Wiz ig Sot Ag eM ho Eve ee Be Oe oe als ok eeelek 54
NEI en Oe oe A GC elsibtege Mice dolnel Smee 54.5 4g avn dle ain 331, 394, 510
Te Og eo PI dats nsec ace cul eiarae Sm mie wince ea 0 deat 198
SESE CORPO ON) C0 ET (0 |: a na ao 436
ERE PAS FESS (eal AA Te MOA Ws AN | EE RE Se 434
SRRMMUNERLGRE Sire csc cAI a Sabina a, oie ahaa, maemo plerete ig at 436
euiraves, method Of attacking. 3.0). 2) fees cn ot bis on Wewleie oe lees 437
EE at Ear ae ar Pa go 375
TRIN aos oo.4-< = Glass. tc oss Por ed x WN rae ahi sour ey re kbar: {mal Stk Oo 8 375
RIEL os Sag ee te ET SIAL cad wi dhk aia Vie wha ath ayes oR Oe ONS « 8 hevara piel « ath 306
EE SCRIE seed Sa hie ch Ib edeie cau. k Fyoa he Ue REE Os ase HE iohie Alea ties 301
Todime; detection in non-electrolytes. . 2... ek ee cet en ene’ 304
FRCRCNOA OE OUNING Coa) oka Us. wks oo Meda avon eha en 306
equilibrium between two solvents. ............. cece eee cece e neces 17
I eg De are en arabe dg ca 'g & oh'vi ge amie Mie wise a, 256. 'g) blaine ow RW RT EER 304
ME OE Ge gre an ys Ge oko 8S ao meee Hees, Lak o> been bight aA AS 305
I crak tL Ge Ie rg tice a Vig bales, oo Rie stile a'n « pbaige a SRI w Wages 4, 10
EM Re Sad Ce ahe oak wa SAG oka Saleh ee oe chKe RM OREO Oe ey a
Nh OOo ida hak ems DAT Ga ee eo ee sll oeapret es eaab 24
concentration proportional to conductivity................. 00. e eee eee 5
TINY So ane oan aod al cuit cranes sea Rie sein of RE poy Aree Ses 5
a ea eco d- ty SSS 5k ald ye ne Yee a 8c oid oa febnc Gua Rialems 10
ROO io or A 27's oc ce CTE HEME Rais Coo ON os ote wewiew oN 4
ee Ua ik yo ASG Tacs 2 vate ork kia we S Sean Ole tMtenen 26
Ionization. See Electrolytic dissociation.
a EIS SS eS eo Ni Nt ara Ie Fe ete we 53
A OURNGOS Ii CONCENLTALION. 6.5 bc... cee cu cee nw sevepese ten 18
SRMMMNE Sg AN Loe Crag MN gayi di ud Bete © ¢ o.c awd Ch SAS Wem eo we 19
ITER OUENN eo gunners LIC Riad Spe hal sv ou di doa vole pape einceda ia aiaolehh aes 8

a
528 INDEX a
‘ PAGE F
Ionization of salts of weak acids or bases. ............ cesccccccccacccceses oa
ei ee Oy Th mee APR MEN RMN OME EG 512
EERO 85 Sins See's a a whe Wey 0k bbe lc GOW pe ke ates Te 504° <5
pL OE OE Pe Pe ae eI merit eae PES Paha ne ST ~142°-49
detection in the presence of ferrous oxide... ...........00ceceeeeeceee , 1488
FUSTI Oi oon Fa KE sie aco en ee pnw Ol piace np ate 143 >
separation from other members of Group IIT.................. 189, 192, 443
state of oxidation in original substance: .............c.ceceeccecececun 4438
J “%
SVE EL OR EP EL PR ETE EET ORT PE EET ELE EERE es 413 a
K 3
Decca. s dal sos pao s 9 ow ca Me pwns vin aS Dye, ubcetdl asi ieee. Wig ee wits ona a en 125
PGGATEDOENUG «065 62d as kvg 5 Soa-s bie oie SoRis,a Roe a, 8-Aie eptb ine Mpcdie win alesktes 5 lee tc 242,
BOUOT TU 26 a5 aaah 0 0 2h-0 no Bis sk lv b's o4'0 9 lass 5 EIST carte ebb ca 93
REMAADEOUS 5's 70s Ss eo Wise are ve Be wee wale vio Nie b Se Pak ae ee » Sheen or
PERSMNIGE G2. hes ce ee S USED ROS WARK TO Seat Poles ¢ ob. eee en 499
L :
Laboratory reagents... os i: 6s s's is 5c ond oes 5% en cisely + fa ene ee ee ag
CES ie 5 6s Scan G4, Vs bn eonTA p/Ble ete bila se, 8p mle thx 9 a altos os bane 55
Pathan 50 ak Fe as oe 2 5 Sree eo ko hE ae pee ke ee 469, 472
ARTA oo Ea keg soon 8 ie 6 le ceo are led nc ge “2 132
Lead, acetate (sugar of lead)... ce ea ee We es 0 ee ee 341
POBCUONG 53s boas vi ae Sk ok oe ae ae FG eR Pe wl ate 205
separation from-Ag-and Hg"... ws. 2 .oicss vss » eoiee eee ee 282, 441
ammonium sulfide group......2......5.0e2cceeeeee 4. 272
: other metals of H2S group................ ee eS 272,440
Buiate;, SOMMION Gio. a ios s a seeds bab nus.. uae dn yl eee ee 209, 437 4
BS Oo GT GG Os MOREE ME OSE URE CAMRY 496
DMG Fo oars OSes eka BRL ENGR O56 ee a oe ee 457, 458
ME BANCOTROIILG 2 ists 5 6c 0 5 8% ‘oa ote eee cle 0's He lee ae wees | Sie ik ne 461
RII 55 ane le wg ale oa tn hl G "oa ohm 6h aa lv'w bs nn. °0 aco se) 9 gk /bo ed ea ghee pe 154
Oa a dR Oe LAAT 102
TS TES Need oe A Pa SR eA a Ne ‘Sete 142
MANS, ATG YES OE oii orace ss tals a tecorae lat shame h prs. soe NEE ale me) 452.
PRRINMEMEG 525-5. ou Sneed Gos ce ve o-b btalarb ig 'asp/asueophasevaehtnvesesk a: Osea eleiw coleee eaten nana 206
BAPTMED eee a os gbei arn wher a bletreal anda! wove SOUL Gao. aoe Sie eet ons lon fig a are ae 458
separation from rubidium and ceesium.............0e cece eee tenes 459
RUINS oan! 6: 6 Worn e so whee a: 0"0' wha wig" tw’ elsb'a Sovoret nse ls bia” 5 Tate es kal mi an ate eye 55
RAMTE OIC aos \ 5's alba! no herarwly’ ole 4 bs ere i Sve-le' where ws aapily a pluie. Shee tsar aaie a np en 224
M
Bineriesin Comenit |. 5 i: o. ioc v.sv sic o's’ 0 vreln.e'e view bh wage. vei elalstele ee q:9seid 5 dle reen ann 94
PUUEKUANTG o's fas 30160 10 4S +o bis aces a Noy Bat ad Bk oom eee 73, 378
RPRMURGLS soos KS 4 harem Aowe Sis, nights a Sybele ee alae Se ae 93, 102
Magnesium ammonium phosphate, solubility of ...........-.eceeeeeseees 75, 95
hydroxide and ammonia... sccses cvs sueecoe€s seb oe seals ane 94

PAGE

TT RIAD Foe ioe adits Coe ess sides ise ve aeceassdwcarebevewe 93

MRRINTENS WP UONE e712) cari ghes hated aos 0 'A/v a aaa ale a'a'e ove 0-0! 60h) ofp a niain oluhy 67

MMPOLIONL IVORY GIRGIICR. 555. sa sac sce siclee 8 ad ewes eee asy > 96, 440

Ns ye Pw ag calsh ¢ cd's ene COM NE alle t viud ou g bene Vers 93, 142

ee Net ae Sek Sy. lic gla ale 6 al vate OU Ria Wek da Ae NO es ORS 214

ee oda foc oe ON 6 dihc sb OSs avy so ee Sade vieu vCadvae 160

separation from Ni, Co, and Zn..................20082 - 189, 192, 443

SE aR oO a SP a ARG 169

eee ON OS ae Nora age aC Ao. 0% 4 dolce bide ¢ oulaye s vie bois 160

ED Foxe OEE Pagar Pat galwe Ee Oe eae eld wee demem giv anels 162

Tg c Sra a 5 595) 4 eg 4 as ak k dog 42 eed eed hb Be 2 vole et 162, 163

OEE EI rg Op nae ak aa ara e gare Bae ae a er 163

tes Se eee U cig see he Gaede eke Uh as web weed Oey ees 102

Re ee A ai! sw eal bea oN ep oe FA eh oe Meee Walbwe pap dee 142

SET ELE LOL ite TER CY RET EEERCC Ee 232

EN eis ine oe oi a. o's ON inay s)njalk o bik's aoe ee end ate es eee 13

applied to hydrolysis of bismuth solutions.................. 211

ORGAO ANG YOAUCHION. 70-5. 6 fs 6s Hs oe clos o eaate 44

precipitation by ammonia.................+. 19, 94

hydrogen sulfide.......... 47, 145, 187

ee SSS Cs oa ATM Ce ae RE TOON 93

EIRP TIE TSS ch ane RPO PSR len er 461

Peeeereeann te OL COFLAIN MOEtAIS 6. je as bee eee cee renee ee ome 63

OD LT 1. SE a ee GPa 157

IIE rah, Sr ck Is Cad sa iw wid we ie Wwe Vie vee Oh 196

RE EE SSS ne eg en eel ne nn a 314

mercury, separation from other metals precipitated by H.S..... 272, 440

ammonium sulfide group............... ¢. 212

TIENT chy ke hah ny oats eT) As aia saab 201

mercury, separation from Ag and Ph............6.. 2.005000 282, 441

1 hy Ge a a ae el cook 5 w bane cw eb oe gate ape 195

MIM AEB rey Ay cise hie eee RA RG Vek ck WRN ae 203

PU UOTIAOMRISEANE EN 2 5 oar a> bce. ok: vilenaite’ ss hnale sesuce dw 09.0 ds Gan tein wae 204

WIGGREMG ANG OXICIC INCTUSALIONS.. . ees ects vee ed cvevessceccseeece 68, 427

Metalloids. See Anions.

a EN EMIT INGO WTOUDE: >. 5 505 osc oe ble ee seb ucee ee ssnscdeameceveges 70

MR re ie ER i Sy awh og a Oe wee walnaiaas 452

general tables for the examination of.................cceceeveceees 440

preparation of solution for analysis for. ...............-0+.000 430, 438

ES ge ae Sieve te ne SEE AER TNE LR re tay: 77

NA eC eee NA a es alesucea Walp ein es 33, 67

Pen ein MMMM BOIUDINIGY (OE... 5 cack cs ce tues sence dee eee eee pescs ees 372

a AEN, DP SLs case wi'sracg 4-0 Fed's Baw vA Re Oe 372

na sarees pen weegidbieale 257, 435, 437

ene Se isos ped ee bene eedoss oy 54, 284

nist Ye foe ats ere Ny aM ces 2c oie 5s die 3.4 dae w a'bb tate wae oe 336

RO ie EL rns vies bad «wate hanes Lane 78, 125
Microcosmic salt. See Phosphorus, salt of.

Mass, COUBCIM ME APOOIIG We ho ss a is hikaleie Gib hs chaboed wcbicedveeeaees 238

BET ON eis Ty oe) ee RIND, sigue ES RLS, i a ee eae ae a 172

530 INDEX
BUNT GONE G0 Cavite gacdis Bs Heer Rore eit PEO vat TH rs va. 6S
DR RGIGD, 555 oj So ac we a's 10 eee ERY Kd ane b aeons hae cele ene 102, 205,
PORTING 0 aay, ib 40:0 jn wees in 6 edbellg clu lag baaoe an OI: ea ee te ei te a x
Mitscherlich test for white phosphorus... ..............cc0cceccees wevevees O
Mobiybdate'of ammonium. 212545505 (4. iis SS Wibieen tale a pale 72, 231, 379 Z
Molybdenum, reactions.............. smi alls‘e/ ace one wo wae Ra oe ae . 491
separation from As, Sb, and Sns. ss... a vo eww lee oe 4 493,
Rintobdenite! 20.0 Me a ee a 491 -
Mbly bic Reid seo ac ok eee Sek ne Whee calculi ae . 491
PU DOGIOE x55 5 Rs .c ajevegnssrn sin 0 -w wak Kia se le" Rel 6 aN ao aed a 465, 480
ET 7 eS: Lk SRM MASE ae NERO RS <4 Oo
TPAOUREBTRI TO ooo ood die saad 0 BWiniol gen dalbp a al, wos. bude. ote -atn ai'b, 9a us SSF a a
DUN TEOO WIG se 550. G ce lo. Va ona bie So wane opus HTN oboe loses ly soe 28/8 Sate ok at 78, 125
EGU Bacio <.creinin 0-5 0d B ass. eto sate th op cmd leg rue aaa nea bs Soc alts ey Pe
Molal:toncentration. <2. 306.26. oak am wae > by Waleed a 4/6 Yi a ee hoe
N ae
eo a See Ea Pn POM ETON AORN 261, 499 ©
BYAGIIUG «onus sls 35 ops nice assess UEP DL Avy 0 bbe tyleteels Torok va ile ea
PREIY Fr hs 5 ss < Sp Seth Sats Stl Le eco VER oe eae ee be,
FOAUPNGHIIE, oo eal ca oe the Ti Meaty eta 4 toes» ach gh ane One aie ee
Negative and positive elements. ............ 20.2 cece ee ene eeeees oe
elements, examination for. . 6..0...03 0. 5 eee sa ve a 0) oh !
PRE STINET SF hd)p atk pion ere c SSe ee a a ede ap SOK iS oat ete a 469, 473
Wornst Tommi s. 50. oP aig Pa ee sige v cgs ales Vo a pn ald eg |e
Nessler’s reagent... ........00eceeees 5s er wslevanle aprplg pa 73 e Saibs tase en 73, 90, 198 _
TE aE ENE LO OTE RUT ye 90
Neutralization of acids and bases... ...... 053 66 ccs) Stine be we ab else 6, 615
eabiof. vc ca ah eo ew ebb ek mel aT ra eee eee eat :
OWING Se cag old ¥ SG a wig a we eRe S ed eee pee . 604 |
OE Le Se OPC Ree eC ba Mmr at Meme 172, 224
Nickel, detection in cobalt salts... .........cccecececetpenswenceveces 176, 1774
peelings. 6 Shs soa is 6a Sate k Bis hve 5.6 OR cs ieee eae 172
separation from Mn, Co, and Zn... ...5 2.0... ese c esc e cece 189, 192, 443
BS 5 22 FS RE GR wa Seste’ ow eipal a aNe}d a. tVagac ees apa ge eas. 174
UOMO, 2b nk ais ole eb a aca ln wry 050 ew a jh a Telet con wha ypc tte ane aaa 480 ©
MOON 5 Vic 5.55 aba God coins hale vere aie lle der om Ses ee 480, 482
separation from tantalum... 00.5 6....sseeeee see dete dee gegen
Witrates, solubility Gf).6 6.60. le oes CaN TS Coe Cae Manis ele er 393
Nitric acid, detection in the presence of nitrous acid... .....%'s35 +) See 395,
detection in the presence of HCIO; and HCl................... . 398 4
OXIMMHON Ua OW fs huis dae ov eS Rasy te hash eee ee ee ee 30 a
he of hydrochloric acid... 6... 66444 ss ee see We He hes eR ' W874
Poncbi) 2) 0 6 sid eas a:s00% ceels pedis var eee poke Han ee: Si)
ER PA em i irae 7 eRe te oe), yes Pele alk» «ee .¢ one +00) 439 79
Se Oe ere ea ae 330 a
Nitrogen, detection in organic compounds, ..........0seecce cece eeeeeeneee 316
WNitro-prisside test: oe kev oe oe ove eon 0:44 shinee ese sie me tie emma 60a
MME, Cee eck ela eink cava dae gusd io accatenthien ae 330
Nitrosoplatinic chloride: .)5 0325 4-.j.+ wis <0 sein se die'uin do biae to's os m8 a ing oe 269

: INDEX 531

PAGE
aE SIMON TIRE MRI clots 54. eh aold caida SMSMAE a they amt oie easy ee oaks eae DP astna ack peas 287
EIN «Sees tesa Pee st Oot cee cee éeed en pee eae eaisd a wakek 330
Nitrous acid, detection in the presence of HNO3............. cece eee eee eees 395
ANOS Sie a oe Paes F FTE Trio ORK ODE Ure Oe Reb Metal e Sine 329
MIRROR Selo os oaks ad: rela ae Ce bial gia'e ov ed epeieia.e she unpdaaeees ered 10
eg a tM get e Seale aig ey acy a’ ce als ¢) 40 vg ayiiene sca oka 3
EN Se aI SB os dhe ghee ae a ya winly cob c'y ey ola wes TE eee nate 6
Non-metallic substances, examination Of............6ccceecececceecceceees 423
Normal solutions................ Le eeate alee Chae SERN od vals be lagi weet 9
SS 5 iN ou ov Sw oa SRG y erelt Siu 00 ke bee e's 00-09 p wld Ko Woe 172

TAS 8s ee Coa sesc a 5 ed lo Wipe we oie Vpn Woah at ds Vow aa Vee 0b .b ete waite 1

Es ae hd sss a bee bolas we abe oa ce bes Mh OS a Sig aia! k abot k 93 -
I rere slat | NS Pe oe tn a-ulely © aye ocx) owls wh ouie, wn Soucek Raw BD 496
‘Organic matter, removal of.:.................. 129, 152, 172; 236, 238, 240, 443
“substances, tests for halogens, sulfur, etc.............. 291, 299, 304, 338
IER oa gtd) ans sialyl vc aloe te: wr ploy has stg 's aha d's aie/s wy et re to ee 125
BN a ie iso a's oan: sha’ be Peas acTen ai Rate tae Wigs Gees Nal ana weld al Bee aes Sar | 125
RI eT Ns ear yo Vi k'v oh wig dln Bia) Wiper eb otwW. Ae Wy love’ oi chuya- wa Cade we 224
i ee eel ge Bp Le arial hip re nsw BAe © aiden. wi hw blaisila Wa We #69 bideots 469
eS oe Pr dS gue a yi plank ole Wik. eicle Cielvig as iefae Slubperes bawes 125
IT ht Sas ee pn Soe lead Silay wR SS aged ok wpe ten s+ Maal 50S
NL UF NR ara Vert sso Sab preci ai 3s iw Vee iowa gs Athy eae ae ead otis 504
POOAMMAIIN, TEACTIONS. :. 2.56.8. eee ees Pa D a anwar Tat alcl Bee haw SES aOR Reb heels re 509
EN Se Tene Mise GW Acd ily AiG: o Stald 0%, m:aShS: souces oo, BURGOS, a She eos ntoth 37
EPIMMEEVION OFT MOTIBION «0.554 odo os clic ols wise vals oie a Uo ee alleles Da wa ely ge 361
LEIA SRB Pa a pl Sea A Ieee in, a aneE tie, oh SMIeS ERE 8 360
RIES Mee Sets eA Seer cic cae Aan oe as a aeie es Sk ioe =o oe ONS 360
detection in Saelinbadiey Ee UNOTIMAL ROE Fs Suh aeradep s aighahss ois oe 424, 443
SR MAAGIOT DY PHOETRATAN IE. 6S sass cv ols bcs alee we so cl biatecb ace 0 Grape Oe 171
Oxidation, and reduction in vitreous fluxes....... 00.0... 00 cece eee ee cece 64, 425
PERNT REDL a ime ga Lo tASS (ok alc ga’ aly pia eas deste pede cae 32, 138
halogens........ Tee RD oh AIAN ct a ALP LEAT OBOE © Fr Be 30
G0 iS kee eee 32
NUN EMRE csc 9) Sata< ert iahe cna (able Vives kd widcbce Be cause Wea bata era rae 30
MPUMNMIPRIATIS TROUT 22h 5705S. SIG oc DA Ras a Bee wTb ad 8.8 BR wie ee ale 31, 171
OE AEE RS SLE NS Ss Sa AE aa A a a ORE CED ASE gy AL gt ae 27
of one molecule at expense of another similar one................ 170
EE DNCER EN aes Tee. 6.55 Sabon, Wie Sun Pacer ws om Riks ANA aba» Weetie Uie ONRTaa ON 287
PIOMNMUB si lds vie ag sie kheee Yee oe ARE AER Tee 286
SER hE res Aa Ske ats ay hla eo aan aceera Ties sa aah 37, 40, 48
gn oy ie Ai chal tow cocoa Gao civ ¥ Bcwlamdre'n branes 68, 427
I Cetra ed ee SOS NE wie g hikalalcis baph@awis PAICEMIC Es 86

Me OSE ERS BO ai OR ny CA a a be WER ee ee 506
Pea EI re ee Sr Re eee Ly vise «led gecslb ala aves dwibiwie gw ale abe 504
Pm RUNNERS Fae 6 oc iy Soe oa ww bap awed (ae bake ewe 505

532 INDEX

Palladous ‘compounds sss... 65 6 fave wiess de deen eas NaW@e ba sagagie v0.9 ose >> ae
PRCA Weis os oo bo wie cites fake owes be Sot Ca ea a Be
PRPOATOORIO. BOI, © co sro: 6 ovo ooo Fe Nery widgets Wh? cote. a 3 eee .

' Perchloric acid........... AN ee eee PP rae . 406
Papen romnd aid so. sa5 0's os vo chidgs s Sok weeding ck a vee 132, 139
Permanganates. See Permanganic acid. . . =
Perea APmasine: MCMC Sos vam oie 4 core to ed cock tw kato a ae ee ee
oxidation DY oy 6.25. 5.55 sa we elute Oe oleate eee 31, 171

Pepe iRie BOI. Sie s 5c ok oa ne eee ee ess ip hekua soe a A
PE MIONG ONE 6S Se hea etdig cab's Sidecar ges ie 488,

a OSes a Rag kak oe Tia ees we Den ee 461
Phewobpithedeha sla seals ae Secs wehbe eth tebe tuk) Laenel oe he
Prorohates; Golubility:-Of. os a sy in Fh y seve sabe ook Pe slo hae ce :
Pilospnides. aonition Of 4 o< i sissies Des ah rag Res woe ee wt 3s 42
PUES fa nna eure hin uisien Sah Neo ac Ree 369
Phosphoric acid, detection in preliminary examination..................... . 443

RAPER. fe, c's. 3's 5.0 SAGA rhe pate e Dual os Bante er oe 382
detection according to Blondlot-Dusart...................0000- 88. j x
Mitacherliels)..c. crs oP can at gina ea 7
in iron and steel. 6.0... 504 sats Le ea eee ee .
pentachloride, dissociation of ........ Pic oar b g wakes See oh ane ari
1 IE Ee Ne ARON Te iy TL ee 64, 380, 425
Physical behavior of solutions.. ....05. 6.02000 4.200 iS ohess nd bee ‘a
TENG Bee 58 hts ha ten ots BOP PG OY ed Se a ee pe
Platinic chloride. See Chlorplatinic acid. 4
ANNIE, TOOUONG. 6 ove Bock gn tow Cohn oid te emily Ninel alate ere RRR ee, 514, 516 7
MONCUIONS 5 °5 35. 650 o's so eee paad wes wal Bas kcle ON ABs te 3
sesidues, treatment Of) 006 5 50'S bv avs Hee
goparation from: BOlGs 5. u65.4 5650 c esse ye bak aed Mos ON Ey
APPS Ace ret i ore oar ee + Lececleds od Waa 6, catey te a
IMMUNE AOA, oss. bas esc vie co Weeks Rew © Lalens © suk soos Geer een

WE cs ko ose vc 0 ook bin, OWa bE wise Eee we ha ed ee Pe ee 156
Positaye snd negative elements. ....5 <0. 6 ess sna be ees 0b nes awte ss 2 eee
Potash PER 6 sho. 6 w'6 sign orb LU aA wre Wi.a 'o. « BRw Ano WSsgee Wekzt ol catwiie ta Ue 2

INDEX 533

PAGE

POtabalM PersUliate. 2... 25. gcc ea ecesewcerecccccvas asdeéceVewsccesear GUE
pyroantimonate............. Pee terior TOPE FP Pee 5 , (
SURMMTITRLON SUMOTL WILN es: bic seis eaics b.t.0.bsb cla aen'ae ve obs 130, 157, 438.

IN eR OR ea abs CRE RE OO SOC NEOR a ehon a ee alee ee coer 78

RCINMAES EEE OS oe, 5 ara’ Wal gad slateie vise id * caleeib hte ha tle eles 67
RMN 8 seo a ig andl ol woth a.a°X «6 SNéip. 410 Gace be 01 Depa 37, 40, 43
Es eet g Nae acu a(n ais e.ar'e sseie ast allele Sinca asedie «4 Mea cie Bw e 491, 494
oe 0 a ea i See re Ae marae A Pa 469, 473
NINN OE G92 eg od ahaa vs gk a's, W ofa ls aS ood Gedo a Mid Na din’s, Mace’ epee 58
TT ORDLIMTALION «, v5). 0). 6, wv cpdsgied h odiels ae sa des co cieae Sddee a 423, 428
ER nd ata sale GG cha een le Midis! eieyein Vosss-n) eA qs ste. a's oud ais bmw Rad: 0e 1
ONT S 6 oo 5. inca 9) dia nik. hic.ere, sep nie, ec gersie'e eas Pde hate eo ae ty (eee 60
Es Seas asco pe en Fas Wie daa FS ae Nl 8 pe IT whee hike ni od ele aie Qinvee's 224, 278
ee orn eS ay ao g bk wae ales ele ose so tan, Led ee ee tate 150
I SRT WOMOIW 6 oS So ohn a 04 210s Geo es Oe Wee islae ole naw walbeete 318
0 SRE ges ie CADE UR RPO Be Meer coh eae Dotan aa PE Pt 321

TIED er aed opt S als Ou. b eo oS aleve edearhi as advan Ne cekiods'e walk 310
RRS ery See EF sca City enc Fda sate wie Ok ovlele nee veel be 263, 265
aa ite ey ayn aisle suai ed ww aN oe wus ck ceeiee Rs pee eS 242, 278
I ee eS wigs hii s aio os sine Ge 'els aceed dyes Oe MMs pia eca ar 142, 336
LS ol sit ig ch cre ah as ee slSycleis Wael ea cate es ee Kwa eee omneh oes 465
EE edo co 5 Gi > fae aieie, cree ae, ¢ WLS owe sin Shea. ace Pele Riadeeew ao ogc 160
RESO a (OR ea ik on Reg ae veal) Lhe ped Rome ‘.. 102, 205
RI SIPIBILUY OD 5c oo. aaa ok a0 she os, Fan aa 9.0 nye" PM Be Rie -eieigt oss De 374
SC MNCINIO’QCIN, TORCUONS . ... 34.0 cies cs Sc bce ecru Sn ee oa Selec oiee ovis ye bie 374
Pyrosulfate of potassium, fusion with. ................ cece ewww eee 130, 157, 4388
Pyroxene-amphibole group of minerals. ........... 2. cece cece e eee e ee eeees 93

Q
TIMED. oa a) ong. 5s 55 cay, so ooaes & vehi os Mia news BA Rho DE we aiale eg ibe eats 1
ENS MATRUUNIS Seo aks os alg lane sie aid a Beard ee dra bo eee Wee: Wo ws wah a ssbrald tree’ é 1
Ie Sete alot 5a F Siatele gin cca ih ¥ ied amie w'o:PeTatg BIR a tue Rien « Hivigrace eels 413
| R |

LOT NE ook. soc vee Put Glen, bw oes leo dea w ON 461, 472, 474
a ARE el SE oea a is or a Sei) | ag cae nee wy Seam ANSE AN eel Ra 456
MEMS ICONS tik et POR Seo ores ba ose Bead ewes ob es 490
ERR ORES OR See MCSE AR TA i alae SR Re 1
EE rte as 5 Poe SA o ey dS. ¢ dew EUs ww ee Pele 26, 51

Rey IEE Se Ae ee a WE Sr eae eat oon SO Re ees OP nae oe Se 26
IE IC SCO 8g fice Fc lg ain Bs bie kus ie a vieie Rely Coie athe nies 1, 61
NNN ciclo wtacalas eke FINES CNS 6 Vio F Rip OTN. Whew Oak ple Ghd Sales 1

Se NERMNC ANG a og le es aer te taty np WEG a8 PLGRS a PEE NIEE OR RS Laie alee 75

Mee TOF CUAILALIVO GNALYBIG: 33.5 055 0 rs W esi cele ca bee ha ane de 1

RIN ots). yeh cr cca g Rt pla Ware yl agteievale<g « vas Ue ee es bile aps Selene aa 1
er IMNTICEACIOU OE 5.14. 255 Vc Cals PAs pose eb ins FON MeE Saat w ews whieh ues 71
ICES 2g a8 s GNOME, Ging SA BR Rte mah db Rat ce ores Whee eile Nie wie es 224
NRE kr LS Ue cats vow hes > VET ss oN Ree gO SINER CERO SS ae sc 207

534 INDEX

Redueton by Devan alloy... FOE: Ta albe La gnd aoe eke vent
hydriodie acid.. sev bee cubes be wh gleae po leery ie se
Rydtegen*paroxiie: ets Sich pred SME Ty PRR ete ble I eee ody an

sulfide....... <i pclae pints Phase Ola! sectal'e HRS LW URS 4k Reon
metala syria: S250} kb en MIA A WERE Pee yr TIS er
nascent hydrogens +5) estes s hc ces vey ele ee eee Beg
stannous chloride..... Se TRU Te eat bas OAS Gal oo vs eh ae
SUN UGG BOG LSS sass vs Seclew Seen ae oe ahs vevsteebl ss) OOpm———

in a glass tube....... wh ATs edt yo ae eon a f
upper reducing flame...................0.68. ae e's pc ca ema a

om chareonla gk. voces Nak es can int yews eg ee eee oe Oe
charcoal stick....... Dingle ath heels tote OR ces Wee . é ain 0

with sodium, potassium, and magnesium.................. ‘au

Reinsch arsenic test...... Cah hechlasy Naa. ey pare ove ee
PRS VOESHMG FOACTIONN, 6.5.55) hie oss = 510 tno ole ig a in 60's wail One +966 en
Rhodium. .... bP S SURAT SRG Suhel kean ee at See ee ame +i, oo

Rinmann’s green. . ee eet y Te eet as Us ass he tee ree ee oa ot
pietie wats 0265.0 <3 c3200: eis eet MPN eat ea iS oa
PKI ACS. oo, PLRe ei. Ghote aL E Uk ee aen easy, heen . eee
sah VORUEONE: i205 sig 84 5, ald is heals wiehn e uloup the Peale sha.» Oe ve ee hee
My & separation. from lithium: and: cesium: 03.5 5..0.5 va sapeie eee ae
Rub Si Ta 5 ye Rie bales Bk o WMD CS We NEE De 2 1k
Peete AGO 6 oa. ae Ra CRE A Ka) Rnd R eee eee ee ee
MEARE GTLAPENTY THACEIONIB. os 56's vg ote so, ce piglets vim boo donk ls retails 2 iat eR wo. oe
MA TAs kha. é ole sed sk aS teh caba ayes eee tas ae Dec Gere ..» 157, 1605

.
- ‘

PE CMMISORIOVING 5 Sia. Cue Os yeas dena ake SUNT E a $ ple einstellen 6 alana
Salt of phosphorus beads............ As Ee We hanglegig an a eee ae ....- 64, 380, 425)
solutions, concentrations of................. da ba RR oe WN 00 ob oot a

PRM s 4s acs 0 pig8 OSs Oe ang a ice PAE ha eeld che ia ea es cca Ba
A MMRAPONU EID 000 55 5-2!2 sio:a Sv ovo ves Kw ecote kt how oom 0.9.8 10Ny 8 = 0 ld ore ek
RRA TAOU OL 5 i:k vb nd sss EC OR Oe SAW TS Ca ee oie glee ene
solubility table of......... Pst Pl Shes VHeeee ee Ga dione sv SIE 5.9 cae
RIE cho cc w deine sc ews wid ean Sd Fede dedee Lube ogee ee oe Sea ~ as LORS
<a aa rape ire nner eNerrarmeny Pirate y reser eS) wih F i oey ene
RMUTHEOD GOMITIONG. 6. bis 6 diel ide de dedave 46 daithcle FOBT) Ne EV Are J a
RN hse) ano> soo Aa.0\l 2 dbut Ad mod. cnt eomn weet Bei ss.atoaktea (1.
Selenic SD; TOR OGIODS oa ol.ol's Wasson ctw ey mp tiieivin dl clapath eae Orne a
Selenious acid, reactions. Cee Wiad tei aie ed CALA eS YBa a” eer.
MIL iris cicsa.+ « lee Seas ere Wrculct sep iee Chee Ladenen bags tn
iid tAtbarinints $a ope eel ean sud ed-caelad vlc ater 1 DS OL
test for in sulfuric acid........ Ld inte Meee ek EER
EE POETS TLIO IS TF Ly PP hikideneniad ee
Sensitiveness of reactions........... daw eatd lidvus.d ee Ree PS ee
Separation of the acids . 6 oie. pais vee cating day ewes oe vee ee eit etal ales 446, 4 49
COTitS MCCAINS) 630s avsivi nel slcw ee ales grrr te eeeeeerees &
five metal groups from one another..............++00+000+

“ . PAGE
Separation of gold from platinum... ........-.:e cece eee cece eee eee e nes 270
MCR RATOUN Ee cua ak cs Chea eet Cra Aes ess eats he 282, 441
GROUD LES AE oss Ne iia aise Se Pe ee: 272, 441
eT LN Ae ORE AES ger Bai beret h aay Ba 189, 192, 443, 444
RRIRTE IN yo 'a sian é cakes wie es eww descr h Se ake 111, 113, 440
COI Ms feo re Sete tay tag hel Mane aia Sie Os 97, 99, 441
MATT APIEIIAY THOUAIS 57.5 5 oe cele 4 i nnece Co dine vee Ou oe niee baa lo’ 514
coy Phi oh 5 pled lore wae a'e aia e ye Mad dal tae and 2 y tege e 93
TOO Cera el yon acd ig Oly ain alaluratain: eet bos ig We > bie Sch Bb. ib Os 102, 142
EE 2 SRS SE ear eae evel stake glove late’e “aed alas aw eyed oe aie 415
ERM EMLT ACID EIS LUROR AT. 6 Gis gicae kote clus 4s foea ceed sve e ep ba es 417
WOFRIY ER es es oii eee Ca ta a eae a the Vado ale Bh 419
I EEO Ste IY) Pee wd aetna diac Draka ald Belt ms. eS wd horn. a0 8 416
MEOEEIC IIMA D IG IXY BOOS oso es o's ce Sials vid ad atele Dara gtece wee tu © brawls 417
EP CIE 5), ahh hatec wits lel aa Bartle Cae aatalee al DEA Ged wee 415
pUOTAGIC’. 6... Ss. 5 AES OPE Sh UNM as TBs PE Rese Si YR Eee ea a ae 413
I a Sash cn Raced fi a lh conteio dle’? acetate; Abie = "9'e-9je a te 8 os 420
emmmner ita 18) OR ONG GUCCI: Soc. cies Wiha ie 3s oa whe Seay bo a bas ghele's 420
Silver chloride, solution of.............. Pee SEN aM er RATE CW -cke oly: el RD 20
IMT Tht ek CL ise GMa su hes Tac he aie a tee bp MgiNeS dpe 028,850 e's 24
hcg fering G:C pines Lee GBR ALOR sb slp CHER awe Oba pint 449
re GUA ens torn bgt tN os ypu a eA oN goin ws 278
separation from lead and mercury... ......... 6.00.0 cece eee eeees 282, 441
MI STC ai gta Weng She Oren ees hua y GATE Sand <ihcd boas Re STR ee Bae <a 499
I ke OT SRT ae 2 aad en ek anew bile divas 6 C4 419, 426
I tects So a 8 pk gia inka 'g ralnrge hie tM acdc s prstew RE RE es ee eee 178
RE es gt nce SK era acc Goh asa uei des bie Cece wad LE wwe bs 178, 224
IIE es ie ce te eat od ile eg ea pe week 6 Wotan aie a ees <aieik y 102, 184
Ne ar, fe ee ee See Oe lS gs he ae Re ete 82
Sodium, carbonate, ionization of.................00.2 eee ae Dette onus 9
cobaltic nitrite, reagent............. WOES N es % x ACL RA Ee 74, 81
RUE OG NG NUE y cde tg hse nN Gea ASS Bim ako. CIEE EAS OR lone 336
ig RO BEIM Sr GE PRS. EARP he Ins Paste Gd, 2 Os Pa eM ON 83
EI SA Got a SPAN Lice iak Reh bp Meets pate: ARR SE 82
I MCMENUME EE, OFA ve at))s SG Clalsgl ved Va cu alayar'gl Mis, wkd pha tage Wk ewe ohe bw Ate 67
IO ce ne eG bie be PS ECR eye Sede Hens RH e weet dye we 8 220
merapautics and. solubility products... 60.05. 06. cc ee wee ee sbincecens 19, 23
SOHO VOR Ee icv oe Ora ba eo Ae Re We oe lacneere Cas 432 .
PmErGe TOM -MNObALIC SUDStANICES: ss... ede.c cose s tenes ccvs sis caccevuccese 423
PUMEMNIIMOE AMEX SUDSTBNCO.: 6s crn osc ccs ceeds dese sdiscchevtdeseects 430, 488, 447
RETA MULTI UIUOEE oa 55. bovy ei ealciey were ee cain ene Hoon 435, 4386
ET on oo 4 vv 55 8 Re Kos + 00 ie nsec iol Wubi 5 4 be Be oo keen ae ee 37.
PERM a rhe Lice 8 s-voncinin's acy wb ip hast paren ach alee arae es 38
NE Sig Sa ut hare sabres cot iia outed Flee EOS Rede Cb eevee ue 452
chemical, electrical and physical behavior of...............00000 3
EERIE A SS ite 9) cepa, Nina @ eRe Sean igre hee car ey ali 58
saturated and super-saturated.... 0.2260. cece cece cece ee uceeene 12, 37
re Re Pt 0.0, eke IL wos We gates ae OA wn’ ov vw pn bi iele Eo a Wa leak aioe 1
I IE Nt rk hg Saree yd wf Xe A ERE 3 eR OA EMIT SL eam 120

Special reagents......... Oe EMT NRSC | AR ANE Ah ECE tre ER a eT 73

536 INDEX
Spaptnogranh. Saas disiv eps th Chia Cease eaters ok a
PHMMTROUIOUOR ge Sad wed vin whee 8 > dhs detent Ale a oR eats eat eee a
Spectroscope.......... "Sos Relea oe a 5 ure Se eee a
Spectrogeapio analysis... Soy. oe os a ee See ee ee a |
Spectrum, measuring the lines of............... Aa 4-9 Gt eee os ia um '
SpMAlATiLO heck eA le cok 244d Ga LR eae ae ee ne ee ia
Spinel... kee ee eee eta eee ova nb Sea a Sead. are teat iy ae ”
Bpadwmens .. 25.5 spose Nadas ces eto os ecb es aekeaeaee ibaa ee . 4
FSCRUETIAD OID oS Merde 6 vince os Whe cee ehow Se ktee ee 5 eee Y
ORMPIOUIGR 25 ois = og oaks wee o's bd Kips See wane RIA ga ee 25
b-Stannic compounds, . 5.62 (22)... 5 boas sce nce hee 0k ben bob glpse lene
Stannous chloride, reduction with........... ida ne tae ee 2
DORNDOWLEIE : 2iits5 soe Paes ts dae os Pea er 251
SGGTINYL CHINTIGS 3.025 ik SSG eye bi ee be atk ee oka + cow eueel peg eee ea . 257
Starch and potassium iodide reagent. .............c cnc cevscecpeveceushes ot
Steatife............9 ELEM Pie ee ei SPER PEO TT Le ee
CLT SS es Seer ee ee ere Rema Me PERN ECGS WS
SME oa wick 6 Ba 0 kee RE iis wlan alps a abu agit way Se ee een 205, 494
Sttength of reagents. ... 4. coc. oes cs gee ad cde dude Sued at bwolie Loa ee 7
PRGNRAANUING. <> a8. v'Aas shes aes cw bes poy Nie oR 102, 106
ieontine, POMOC he lad a o-ddiain set ene Se bs ADIN ie loon ae ..
deparation from, Ba and Ca. i 6.30.6 cee ok oe a ee 110, 440
Gulkates: solubility of....... sae ara.a eal hacbrn,a ws 4 LD Sasa ale cadre 404
Sulfide, detection in presence of sulfite and thiosulfite....................00- 390
Precimitation, thoOry Of. .s.4 9 cee ces ous coin ks caer meee ae 47, 186
Suifides, behavior on ignition... 2... 00.5 Sede ss cess e ce kke ars ao diene 3
mee by Of. 2. ek td So eatin pee 3 5, VP eR ecules in Ae el 335,
Sulfites of alkaline earths, solubility of..................... eek 5 ee 348, 391
Sulfocyanates. See Thiocyanates. a
Sulfocyanic acid. See Thiocyanic acid. a
PME tied on Ss 5s Sum aioe ch oo oh hae, 4a het ak fens a 339
dioxide, preparation of......... GA's Cede Wace aad aap Ae le oe 347
Sulfuretted hydrogen. See Hydrosulfuric acid. mo
Sulfuric acid, ionization of... ........5....0.. 00 wi eed de bss «alone 8, 10
TOAGIIONE oo le eS scarhe Sa ee PE Oa RR ata Nee ae om ae }
tests for selenivm in. ..6 040.000. as Ges ss arkalcd x ieee ee
Sulfurous acid, detection in the presente of HS.03.and HeS .. 2.24 se oss on nee
TOACHONE <x. Ci anc « © Bao ale Balto gaa Bae Bika phn ae
Teduction DY 64. 55 0 34.0 d4 oe be tin u wrolete oath slate pea iets tate
Supermaairated solution... asses asle saws Anbae Oe eg ces bebe
PN a ooo ha 03.0.0 0 0.0 bye Hb Ow 5 See es gE Fa a aE ee 261, 499
RII ns vb.o son te ua se Mink wg oa Ue Cee oe narele a58 6am and amen tine aaa
SSAA OS ey rere hs COTE aa
cL
Ae oad a) su on 'dvin ne bc bb 0h 5 we A eT ceR lye We oie Oactre abe Deen ey
(ee Le Sees SU te es er
MPM FORCLIONG, «6. . 5-5 0.5 EE) Foe € 99 80 bev Re ise kee a ae nr
separation from nicbiim: ) . 3. 266.555 eeu va ose se cee hee ae 48
NO oa oso 8 Tas 5 cha meee wk Re PL Gnane ieee 244, 362

: a

PAGE
NREL GEEE 2 ise ios e's s cee das oo wae Nise Ciclo EAL AT AIOE Tt EES 362
artrates, solubility of..............s.0005- OR CAR ip contd SAT ea GARI eae pak 362
Taste, sense of...... eee gra a Milas dee 416 GR TEAR Ts SWRA TES wars we eee where 2
NOL MTU CoE, sag GN a pets ao heb b-Aie wes Go's! 0 vishe igiaw Fie Ge pb 520
EE Se Pha as Cae Bac gag in git bie pam Bi XSiriaiphin A o-ciy 43 tgs wie GRIP Ro Ie we ® 499
5 RE DRONITIVL SER MUNN 01 ile a0 cae dino hee Gale G lg sew Sie dog w eeine OL oe 502
I SORCUIONIA Ss ganic dos wie lane a. © sted <sih'o wpa wien #0 6 4b os soi B hiece es 501
EDS eat Ai Ne Gf eis Gls, 5's bls caw wk ce 2 np Oe 9, 8 hoe CGO HIN 0 ye BN 486
Thallium, reactions...............-+- fl Cal Pe rae A Aaa BER OREN, Sle SR 485
REIN, 00 ho Le args 2. ola Bb insm ee Rae $Y dpe + Lidia Npleibloe Cave wb alps 485
NE SR ec he Ae Day be sae bhaey sel tieeeece 131
Thermonatrite.......... Preece Kew TER PER aaa iets ie alamo ie to 82
SER ste el gS Ste 8 a gs ig gin nt a acs oes & Klas WEED Nivree 0 tn 323
in presence of ferrocyanide and ferricyanide................ 327
SERIE ID LS CVE ose o olacaks bone $14 dbcte ora sions ue bits Gee bie elelp¥ ve alo as 387
Thiosulfuric acid, in presence of sulfurous acid and hydrogen sulfide.......... 390
RORUTIOU. Lies Cie hee os on cs SN Oe ee te eee 387
Ree a ere ity on Sy aake pa uaa Veneer amass 36 45 465
RES Re TEA Be ln ee a ice on OA Ca Oar 465
Rl Tee tae aon ce Gry Wat eMidia Dx ale Pils gin! oe wets Ee Swchld w sia sis 465
NT ria ko a ay fh secly phere GR Ae VIR UN'S a 'ee So) Soa es Lest WSl ole w ae 496
ce Ry pPre Le re tig chy Wk apg ah WED. Jlabe' ad <i Zip diab @ wis 4g oe 6 249
Separation from copper Zroup.. .... 5.6... s 1c cee cee eseee 272, 275, 441
Ne og ae x Vis hg a 0 c.p ei ace atv ao ype ba daw Se 82, 357
5 eS i bc ienajy, sc e's S's 7s.a.v on SS ay Shae a Riri ee Pe 249, 260, 437, 462
Se Pott ca ig SP cs he oe STA yan, 4:8 ipl aagyid Lies be pie CORN ole «eb be 8 157
Titanium dioxide, analysis of.............. ee avis cst Gish TOR WAP aie Bs ate 438
MMMM EER Mire. AE Peet cse aie Sik se Le ENG « cdipial wiece 157
SS Page oe a Ae wh Ey gk Grae tale 2 aces Wand Sd ME Gh Bee aes des 0 8 125
, ENN Was eee ks ald 4 ie oes Ww s¢ Ceo PIG ans Sass 125
MIPECTISG OL... 55s ccc cee ecw anes LE RIE: cc Bi” PO ey ie EE 2
Lo eh Ay oy eid y coy alice Bow oidclghe Mee e% ¥ 0 vein Saieih's 93
IN eerste gies LS bre dio ie sk a eek s Wald tawe sy, xs MOIS Sarees CLK DOs eRe es 82
CS 5 88 clara Ga tha acgp 90s Ps pe ROA bon bv ee Ud we PP oe alee e 494
ST acre 2 ea iG geo SI wer SONG, Ge in I ha a Me yb Loa lates 494
NI 3565-2 snes en lacgod Sree p'a bls EM ae sad Me OU ah 55, 74, 358, 464
Or ay 2, ab are Bk SIaw Nie kd eee ovine s Boat bNG hae helee eee eee s+ | 147
Re sc oot Sig peleodin wb eee ON es C3 pm ONE ce weld eee Ohne Corea’ 125
U
a oS ios oes 6s. ood Pe da wee sic ces nee abe ew enweh eens 172
oir inc ee ae atc ks Cen asda oe soled Veweehdeuunebpecse 118
Ee Pr kan eens kG Gg indy s ek An ole va py oard@ a reno dueone se 118
a SR URE Ag Rr RT PORE Fw,” eRe et Ret et Oona 154
ihe be eet esata ae aka eresi'o. ew 2 Who WE AGE GOL Bea we 154
Senurevion from Fe, Al, and Cr. . ici. dee ch ve en 189, 443, 445
I ra oe ate nats Ae rg teas «Tete «aR k oe Bein clan oem 154
I IIMRORE 00 Toh ciy ook ec ve eR eae ORR 8 OS 87, 333, 396
NCL WECM PSOTIG IN 2 i S6k ace Fics bev av es ve Ske ss ua e's woe ene cus ees 238

RTE: TCTOURY Bi Oo cd's Fetes Rae's 6 wate a ae e's ve ee adepees 203

5o2

: - { ,
538 - INDEX
oP: os
WMMEIDME IS soy cca eae ee eee ee Myce uape pad sea be Geeeahy 5 a
Vanadie acid..... Lae wire e's uhh Robo aed hicks ides dhe panna des vahe eA apes «el
Vanadinite....... A POE ee are pre A TMi tren AL Say go 102, 205,
OMARION =, dean Wn Viera sip Rebar aie a eictents igs 48s ASHEN MeES EE ET ait mA
. detection ‘Mm: focks 5 56s See os sie ee oe ws ed elec nea
VARA PBIB eo cee ikea Sp Oe aad ace ee ete
WEMERIGO oo c aioio. oie ora pip. sib a sales lovee le Muse ecg aus ole bia Rite ak ha See
Volatele substanoess os. 2s sisi. csyepecaicle: 6b \a!p a 3ladueipte'a aie eo 0a mR a eee end
Volatility testi. esc. e e cere int 80h an emlece dap hieke Tan > Piece Jean
W
Waane. law of minas-aetion. ...)) 0 ico chic e'sa ss,s eae 3 9 1a eee aie
Weall-pavers, ATSONIC AN. 1.6. bo vies pas ep ebeie se 5 05,6 Wale ema eee
Walker’s nomenclature of ions............... Se situ « + scp} 0.0 etn RO vee
Washing precipitates 24.5542 shee eats rae ete * stiecs hw each ave a . eee
Winter “smmonia in: os d,s CoS as s eee evn co eos okey Cee eee 43
free from ammonia... 2300655055 cds cnntensryy atone: See ae
GEIS Sis oiece pH A ale laieg: Oe wha Re o Cs dich baba oe ae 4d
heat of formation from ions.............0....00e00: (ius uae on
MRRPARIOR OE sass et a sleeve eos Waa Ue ey eh wae lle Suge pace
MIN Se hs ae ao SETS SOON ee ae Me att ae ae
SW ereaenvartdae: 2 ooo Se ee os LAS EEK Lh aA cee ode AOMORI Oe Pe ee ee
Wave-length spectroscope. 2... 2.0605... ene bee ee pees one a ens cle eee nn
Weak acids and weak bases, salts of................. eee
Wet way, reactions in the..... gift sigarete csc RG ca 4 b,8 7 Rm aoe ata a -
TL oO ieee rR SPORE A TIMES ARPA CM Nae IN ee RE 2, 10
WOhienthes ss Sa a eee ices» ive Cl Fa ieee ee
WN RING cs sce ocak oe eu bo biel wine web elie ccd aie con ape ip oe le a
rem cc Sy he Ste Ro a ge er 205, 491
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x 2
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RENE 54.605 vib tb aalaud HOPS Ae Ayah ean eb a vee ko iene ope
eameintAMRth:. 35 6 oooh ca doen, tw ce oa es ee 467,
TRUS DOUAAOTIB, 2. 6.5 so 60S ive sa nile ave stale Ghee a 0a wivtelle Site ol ae eae
separation from Mn, Ni, Co, and Zn..................5 189, 192, 443,
PRIMA eas. 5 os, bs iged HGRA bib coe M Ee Ate RIR LL EE ODER Oe eT Re
Mite made, precipitation Of; oo u6 os d es wd se de bel a AF ene Bip el sae’
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